Author Institution: Department of Chemistry, Miami University, Oxfor

Evans, Marna I.

Taylor, Jay E.

The Ohio Journal of Science

KnowledgeBank at OSU

THE MECHANISM OF THE IODINATION OF PHENOLSJAY E. TAYLOR AND MARNA I. EVANSDepartment of Chemistry, Miami University, OxfordThe transition of organic chemistry from an empirical science to a sciencehaving a well-established theoretical foundation has been slow and most incom-plete. One of the main factors which has retarded this advance has been a lackof knowledge of the nature of organic reactions. Many theories have beenproposed, but proofs of these theories are too often lacking or are quite unsatis-factory. This work was therefore planned to give information concerning themechanism of one reaction; namely, the iodination of phenols in aqueous solutions.Appreciable work on the iodination reaction has been done, and the conclusionsconcerning the mechanism of the reaction have been several. It has been proposedthat the I2 molecule, the I3~ ion, the hypoiodous acid molecule, the I+ ion, orthe iodine atom is the substance directly responsible for the substitution of theiodine. With respect to the phenol both the undissociated molecule and thephenolate ion have been suggested as intermediates. It has been further pro-posed by several authors that the reaction may take place by means of acombination of several of these intermediates. For the most part the proofs ofthese proposals have not been well-founded, and in several instances they arehighly conflicting.For this paper there has been made a quantitative study of the effects ofvariation of pH, potassium iodide concentration, and temperature upon theiodination of 2, 4-dichlorophenol. In each individual study all the other factors,including ion concentrations, have been held as nearly constant as possible. Thecompound 2, 4-dichlorophenol was specifically chosen for study since only onehalogen may be readily substituted.DISCUSSIONThe rate of iodination of 2, 4-dichlorophenol varies markedly with pH. Asmay be seen in table 1, the rate increases by a factor of about 42 as the pH isincreased from 6.0 to 8.0. There are several possible explanations which arelisted below.TABLE 1Variation of the calculated rate constants with pH and temperature. Startingcone, of I-i and 2, 4'dichlorophenol = 0.00183 M. The average deviationfor most of the rate values is between ±1% and =*=#%.pH6.06.47.07.48.08.08.0Added KIm/10.0080.0080.0080.0080.0080.0120.008Temp.° C25.625.625.625.625.625.633.5kmoles"1 sec"10.0760.2630.991.873.181.096.61. Assume the phenolate ion is the active intermediate. Then, the variation inrate will depend on the actual concentration of the phenolate ion which in turnwill depend on pH.THE OHIO JOURNAL OF SCIENCE 53(1): 37, January, 1953.38 JAY E. TAYLOR AND MARNA I. EVANS Vol. L I I IThe ionization constant of phenol may be expressed:where Ki = ionization constant, [$0H] and [*0~] are the phenol and phenolateion concentrations and a — x is the sum of the phenol and phenolate ion con-centrations at any given time t.ThenSubstituting in the second order rate equation:where the starting concentrations of I2 and total phenol are the same. Theiodide ion concentration is assumed to remain constant. Tiessens (1931) has giventhe ionization constant for 2, 4-dichlorophenol as 3.1 x 10~8. In like manner,the equilibrium I~ + I2 +^ I3~ will serve to reduce the rate of reaction, and theequation for the correction due to this equilibrium at a constant pH is:where Kd = 1.4 X 10~3 as given by Van Name and Brown (1917).2. Assume hypoiodous acid is the active intermediate. The constant forthe hydrolysis of iodine (Bray and Connolly, 1911) has been given as:Upon substituting the equation for the dissociation of the triiodide ion into theabove equation we obtain:At an effective iodide ion concentration of 0.0065 M and with a — x equal to0.00183 then:There will be a change in the iodide ion concentration from the overall reaction:The above, equations will, however, be reasonably quantitative in accuracy.3. Assume both the phenolate ion and hypoiodous acid are the reactants.Then, the rate of reaction will roughly vary inversely with the square of thehydrogen ion concentration,No. 1 IODINATION OF PHENOLS 39In table 2 there are compared the calculated concentrations of the phenolateion, the calculated concentrations of hypoiodous acid, and the observed reactionrate constants. In figures 1 and 2 these values are graphed. The plot of thephenolate ion concentrations against the observed rate constants is a straightline and indicates a direct proportion between these values. No such relationshipexists between the hypoiodous acM concentrations and the rate constants. There-fore it may be concluded that the reaction takes place as a result of the interactionof the phenolate ion and the iodine molecule.TABLE 2Variation of the reaction rate with pH. A comparison of the calculatedphenolate ion and hypoiodous acid concentrations with the observed rateconstant values. Temperature 25.6°; added KI concentration 0.008 M;starting concentration of h and 2, J^-dichloro phenol = 0.00183 M.PH6,06.47.07.48.0km"1 sec"10.0760.2630.991.873.18Calc.Starting[C6H3C12O-]m/1 x 10s5.513.244.80.138.Calc.Starting[HOI]m/1 x 1081.53.725.37.150.00Xbu(0I100L0uz0uI 4 OI 2 OI O O8 O6 O4 O2 O//////r//- ///r//'/i e oI 2 3RATE: CONSTANTFigure 1I 2 3RATE: CONSTANTFigure 2FIGURE 1. A graph showing the linear relationship of the starting concentration ofthe 2, 4-dichlorophenolate ion in moles per liter and the observed reaction rate constantsin moles"1 sec"1.FIGURE 2. A graph showing the lack of linear relationship between the startinghypoiodous acid concentration in moles per liter and the observed reaction rate constantsin moles"1 sec"1.40 JAY E. TAYLOR AND MARNA I. EVANS Vol. LIUTwo series of experiments were run in which the iodide ion concentrationwas varied. Table 3 gives the observed rate and compares the calculated I2and HOI concentrations as predicted from the previously derived equations.It may be observed that neither equation quantitatively predicts the markeddecrease in rate.The activation energy has been roughly determined. The value from data intable 1 is 17 kcal/mole.TABLE 3Variation of the reaction rate with iodide concentration. A comparison of the calculatedconcentrations of molecular iodine and hypoidous acid with the observed rate constantvalues. Temperature 25.6°; pH 8.0; starting concentration of I2 and, 2, 4-dichlorophenol = 0.00183 M.AddedK Im/10.0080.012Approx.Effective[ I " ]ra/10.00650.0104k31sec"1.18.09Calc.[HOI]m/1x 1071.56.3k //[HOI]x 1062.11.7Calc.[ I . ]m/1x 1043.22.2k /l\XI 2 ]10*1.00.50CONCLUSIONSFrom the data concerning the variation of rate with pH it may be concludedthe reaction of I2 with phenols occurs by the direct reaction of the I2 molecule andthe phenolate ion.The reason for the very marked effect of potassium iodide is not clear. Pos-sibly another as yet undetermined equilibrium is involved. A possible mechanismis the following:H \HThe equilibrium formation of the I~ in the second step would tend to be reversedby the presence of excess iodide ion. This equilibrium in addition to theI3- i ; I~ + I2 equilibrium might explain the marked effect of the variation ofiodide concentrations. This work with varying iodide ion concentrations clearlyindicates that the mechanism of this reaction cannot be completely outlined fromthe data in this paper.The low activation energy of 17 kcal. would seem to preclude equilibria of thenaturela % 2 Ior I2 t; 1+ + I-,since much larger amounts of energy are required to transform the I2 moleculeeither to I atoms or I+I~ ions. This value of 17 kcal/mole undoubtedly doesnot represent the activation energy for the reaction, I2 + $O~ —» products, sinceboth pH and potassium iodide concentration affect the effective concentrations ofthe reactants. The activation energy should decrease with decreasing iodideion concentrations and decrease with increasing pH. (This conclusion is based onunpublished work on the periodate-glycol reaction by Jay E. Taylor and B.Soldano.) Thus the true activation energy should be appreciably lower than thevalue given.No. 1 IODINATION OF PHENOLS 41It is well known that IC1 is an iodinating agent instead of a chlorinating agent.In view of the above evidence the most likely explanation is that the IC1 moleculeis polar (and not ionic) with the I atom having a partial plus charge. Uponattacking the phenolate ion the I is oriented to the ortho or para positions and theCl to the carbon holding the O~ group. The Cl~ is then split out leaving the Isubstituted in the benzene ring.Work is now being continued with this reaction, and it is hoped that a morecomplete view of the mechanism can be given in the near future.EXPERIMENTALAll chemicals used were of C.P. grade except the 2, 4-dichlorophenol whichwas recrystallized with much difficulty from a mixture of benzene and petroleumether. It was necessary to carry out the recrystallization under anhydrousconditions as the phenol tended to absorb moisture from the atmosphere; thisprevented the formation of crystals. It was kept in a dessicator.The buffer solutions were made from varying amounts of hydrochloric acidand a constant amount of dipotassium hydrogen phosphate. The concentrationof K2HPO4 added to each reaction mixture was equal to five times the concentrationof phenol.The 2, 4-dichlorophenol solution was made by weighing the solid.The iodine solutions were standardized against standard sodium thiosulfate.A weight burette was used for greater accuracy and convenience.To make a given run the equivalency of 1.814 ml or KI -f- I2 solution (measuredfrom a micro pipette) was determined. An equal quantity of the iodine solutionwas then placed in a small glass bucket. The bucket was set in a weighing bottle,and the bottle was then put into a constant temperature bath. An equivalentquantity of the phenol solution (about 18 ml) and the prescribed amount of bufferwere placed in a 100 ml glass-stoppered flask which was then set in a constanttemperature bath. After the reactants had come to temperature the iodine wasdropped into the phenol solution and mixed. After a measured amount of time3 ml of 5% phosphoric acid was added to stop the reaction, and the remainingiodine was titrated with standard thiosulfate.REFERENCESTiessens, G. J. 1931. Preparation and properties of the tri-, tetra-, and penta-chlorophenols.Rec. Trav. Chim., 50: 112-30.Van Name, R. G., and W. G. Brown. 1917. Tri-iodide and tribromide equilibria. Amer.Jour. Sci., 44: 105-23.Bray, W. C, and E. L. Connoly. 1911. The hydrolysis of iodine and bromine. A correction.Jour. Amer. Chem. Soc, 33: 1485.

The Mechanism of the Iodination of Phenols

https://kb.osu.edu/bitstream/1811/4001/1/V53N01_037.pdf

The Mechanism of the Iodination of Phenols

Abstract

Similar works

Full text

Available Versions

KnowledgeBank at OSU