133
A THESIS SUBMI'ITED IN CONFORMTTV WlTH THE REQUIREMENTS FOR THE DEOREE OF MASTER OF APPUED SCIENCE
A THESIS SUBMI'ITED IN CONFORMTTV WlTH THE REQUIREMENTS FOR THE DEOREE OF MASTER OF APPUED SCIENCE
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A study is preaented on the propertks of organic iodides. emphasizing the
Henry's constant within the temperature range of 257S°C. The types of
cornpnds studied were alîcyl iodides, aromate iodides, phenolic iodides, and
iodoalcohols. The values masured ranged from 0.6 f 0.1 for iodoheptane at
25*C. to 1.2 x I O 4 I 500 for iodoethanol at 25OC.
A new experimental method invohring a stripping apparatus, and
radioadively tracecl organic iodides was ernployed, and found to be efficient in
the detemination of Henry's constants.
Various propertbs wem detemined for a large set of organic iodides
through the use of mdels. Conœming the modeling of the Henry's constant,
Mo calwlation meaiods w c ~ e examineci. The fimt method analyrecl invohred
equating the Henry's constant to the ratio of aqueow solubility and vapour
pmssure, and then using exœss enthalpies of solution and vaporization to - .
estimate the temperature dependenœ of the Henry's constant. The second
methoâ was the utilkation of structural bond contributions. The first method was
fwnd to be a&rate'for most of the'orgkic iodides tested, while' the second
- was bund to ôe total& diable only within the range of alkyl idides. .
ACKNUWLEDaMEhnr)
I would like ta say thank you to the following people...
Professor Gmg J. Evans for his patience, the excellent guidance given to me during the completion of this project, and for just being a good guy.
K.H. Lin, without whom innovation and contnbuti'on, this project would not have been possible. Thanks Kai.
Professors C.Q. Jia. and D.W. Krk, members of the advisory cornmittee. for their valuable comment8 and insights which contributeci to aie successfui completion of this work.
My colleagues in the nucfear and environmental group for al1 aieir invaluable help.
Finelly, l'II take this oppoihrnity to thank my friends and family for all their support, and for putting up with me in general.
A k b r c t
Acicnowkdgments
Trbh of Contanb
Lkt of Syrnhk
Lkt of Figuno
L M of T 8 b h
1 .O INTRODUCTION
2.0 BACKGROUND
2.1 HENRYS CONSTANT
2.1.1 .1 WC of Molecular lodine (12)
2.1.1.2 IPC of Monatomic lodine (1)
2.1.2 ESTIMATION OF HENRYS CONSTANT
2.1.2.1 Correlation to connecthrity indices 2.1.2.2 Estimatkn using Ouchiml contributions 2.1 3.3 Estimation h a d on aqwous ralubility
. and vapour pmswm
2-13 IPC CONSTANT AS A FUNCTION OF TEMPERATURE
2-1.3.1 Estimation of entnalpy of vapociration 21.3.2 Estimation of €xcesrc Enthalpy of Solution
2-2 VAPOUR PRESSURE
22.1 PRMOUSLY DETERMlNED VAPOUR PRESSURES
2.22 ESTIMATION OF VAPOUR PRESSURE
2.2.3 VAPOUR PRESSURES FOR 80UDS AT STANDARD TEMPERATUTE AND PRESSURE -
- - . - - - 23 AQUEOUS SOLUBIUM
' 23.1 .PRMOUSl;Y: DETERMlNED AQUEOUS ' . . ' SOLUBIUTIES .- . . . . .. . - . . . -. . .. .
23.2 ESTIMATION OF AQUEOUS SOLUBILIM
23.2.1 Dependana of rdubility on solute ske (Theoretical description of solution P-1
2.3.2.2 Correlation b esomated aqueous activity ~08fRcients
2.3.2.3 €stirnation using structural contributions 2.3.2.4 ComlaUon to odanol-water partition
cosf~ent (Ka) 2.3.2.5 Comhüon to connecüvity Indices 2.3.2.6 Estirnated aquwus rolubilit&s for solid8
at standard teinparaturu and pressure 2.3.2.7 Accuncy of estimation mahodr in
comparimn b literatum value8
3.2.1 RADlOACTlVE TRACING OF ORGANIC IODIDE
3.2.2 TRACING FOR 2-10 WETHANOL
3.2.3 TRACING FOR MOLECULAR lODlNE (12)
BATCH AIR STRIPPING QROCESS
ANALYTICAL PROCEDURE . 3.5.1 Dt3ERMtNATlON OF-GAs PHASE
CONCENTRATTON
3.5.2 DETERMINATION OF AQUEOUS PHASE C O ~ E N T R A ~ O N AND THE RESULTING IPC
4.0 RESULts AND DISCUSSiON
4.2.1 EFFECT OF MOLECULAR STRUCTURE
4.21.1 lodoalkanerr and lodtmromatics 4.2.1.2 EtJbd of Hydmxyl Substituent 4.2.1.3 lsand HO1
4.2.2 E f FECT OF TEMPERATURE
4.2.3 EFFECT OF SALlNlTY
4.3 ESTIMATION OF IPC
4.3.1 U(AMINATl0N OF RATIO OF AQUEOUS SOLUBIUTV AND VAPOUR PRESSURE
4.3.2 SELECTION OF MOST ACCURATE MODEL
4.3.3 IMPROVEMENT OF BOND CONTRl6UTlON MODEL
4.3.4 ANALYSIS OF MEWOD TO CALCULATE HENRY'S CONSTANT AS A ÇUNCTION OF TEMPERATURE
4.4 CORRELATIONS INVOLVING THE IPC
4.5 APPLlCATlONnMPLlCATlON OF RESULTS
6.0 CONCLUSIONS
7.0 APPENDICES
EXCESS AH
SOLUTION
lodine Parabion CocrMnt
Enthalpy of Vaporitatlon
Excess Enthalpy of Solution
Zero Order ConnecMy lndex
Zero Order Valence Connectivity lndex
Fimt Order Valence Connectivity in de^
Polaritability lndex
Aqueous Activity C o e W n t
Temperature
Vapour Pressure
OctanoCWater Parblaon Cœffident
Boiling Point
Molar Volum
Aqueous Sokrbility
Concentration
Figure 4.1 :
Fgum 4 2
Figure 4.3:
FQU~W 4.4:
ngum 4.5:
Figure 4.6:
Fqun 4.7:
Figure 4.8:
Figure 4.0:
Cornparison between wperimental and litemtun 15.131 lPCs for ioâomethane and Mine
IPC vs. molar volume fior habgmated organic compounds
Henry's consîant vs. mdsr volum (br oganic compounds
A) phenol fiwm of Clodophenol. 8) en01 brm of 4- lodop~ol
Bshaviout of IPC with respect to temperature
Clausius-Clapeyron test for orgenic Mides
lodobenzene peak; Temperature: 7!X, Tirne: O hours
lodoberuem peak; Temperatura: 75%, Tirne: i hour
Efkt of electrolyte on IPC of fflomethane
F i u n 4.10: Examination of constant bond contribution auumptbn for alkyl Wdes 74
Fgum 4.1 1: Excew cmthalpy of rdution [J/w vs. molar volume [mWmolJ 80
F i u n 4.12: Obsewed mlaüonship batween q ~ u s rdubility (SI moUL) and the IPC 83
Figure 4.13: Obsennd mlaüonshlp be-n vapouf pnuum (Pl Pa) and the IPC 84
Figun 4.14: Obrawd relationship betwm Occariokrater partition cœlllcknt (&,) and the IPC 85
Tabie 2.1: IPC of mettiyl Mue (51: p 2 / ~ ((mm Hg at 2 5 ° ~ ) l ( r n o ~ ~ ) ]
Table 2.2: IPC for idomethane [6]
Ta& 2.3: IPC resulting from ratio of aqueous dubility and vapur pressun for iodofom (CHI,) [8)
Tabîe 24: IPC values meagured at 298 K (91
TaMs 2.5: lPC data for I,
Tabk 26: lPC data for HO1
Table 2.7: Conœnbation eflisct on 1PC of HO1 at 20% pH 9 [2fl
Tabk 3.1: Detemination of IPC (br CHJ at 25% from expdmentsl data
Table 4.1: Cornparison of h s u r e d and Liirshrre IPC8
Table 4.2: AH- for organic iodides
fable 4.3: Examination of Ratio Assumption for IPC
Table 4.4: Compaflcwn at 25% of M ~ ~ u W IPCs, anâ th- calculsteci using the bond contribution method of Meybn and Howard 1341
Tabk 4.5: Dasrmination of impmveâ C d bond contribution
Tabfe 4.6: ûetemination of improved C--1 bond contribution
Tabîe 4.7: Determination of qwnüta.üve CH, contributign b AH UCESS soumon . .
For organic iodides present in the environment, volatilization fiom water
bodies is a signifiant environmental pathway. This volatilization is also of
interest in relation to nuclear reactor accidents. For exampls, the majom of the
airborne Mine within the reactor containment was in an organic fomi following
aie accident at Three Mile Island [Il.
The environmental hazard posed by radioiodine is very dependent on its
chemistry. Any iodine released ftom the fuel is released predominately in the
fonn of non-volatile Io. However, within an aqueous environment, Io can be
oridized into volatile chernical fonm such as molecular iodine (12), and its
dbpmporüonation product, hypoiodous acid (HOI). Readions between trace
organics and these species of iodine can produce volatile organic iodides. As
mentioned by Beahm et al. (21, organic iodides are produced within nuclear
mactor facilitier when iodine reacts with organics present within the containment
structure. 'The largest sources of organic material are cable insulation .and
jacketing . m e r sources of organic materials include paints, seals, gaskets,
connedon, circuit boards, and lubricants. - .
. In order 10 be releawd fmm a nudear facility, either thmugh delibemte
.venting, or through any aperture in üie conteinment structure, radioiodine has to
banme airborne. M e n aiibome,' radioiodine is much more mobile and
- dangeow than -ït is jn the aqueous phase. The main factor detemining the
extent to which radioiodine becomes airbome is the iodkie partition coefficient
(IPC), defined as the dimensionless ratio of the aqueous and airbome iodine
concentrations.
Ulomately, the partitioning behaviour of organic iodides between the air
phase (ex. the atmosphere, soi1 vapour) and the aqueous phase (ex. lakes,
groundwater) will affect the environmental fate of the contaminants, along with
aie performance of in-situ remediations.
In addition to their devance to the nuckar power industry, organic
iodides released from oœans are also believd to contribute to ozone depletion
in the lower stratosphen. Calwlation of the volatilkation rab requires a
knowledge of relevant IPCs, since in the calculation of overall mass transfer
coefficients it is commonly assumed that the solute concentrations immediately
on either side of the air-water interface are in quilibrium.
The objective of this thesis was to expand the sape of existing literatura,
as it relates to organic iodides. The property of primary concern was the IPC.
. Anoaier goal .of the project was to develop and apply a new experimental
method for the detemination of lPCs baseci on the use of radiolabeYed species
in'a stripping method similar to that usecl by Mackay et el. (31. This method was .
. ' extend@ @- themeasurement uf lPCs of inorgank iodine spedes, 1, and HOI, as
the partitioning of these is ako of interest to nuckar safety .
The cbnditions studied varied from ambknt temperatures to 75OC.
-9nerally fot IPCS, data k only ;vailable at ambient temperatures, which b only -
knowledoe of the partitioning behaviour of organic iodides at higher
temperatures is needed, as it relates to the potentially high temperatures
reached in e nucbar reactor, abng with themally enhanced soi1 remediation,
proposed to overcome mass transfer limitations.
The resuits obtained were also utilisecl in the improvement of existing
models for estirnating IPCs. These m d d s had not been previously applied to
organic iodides in any detail. The pmject al80 examined a method for
interpreting the IPC in ternis of excecw enthalpies of solution and enthalpies of
vaporization. for the purpose of calculating the IPC as a function of temperature.
Finally, the detemined lPCs were used to develop armlations, for use in
estimating other environmentally relevant propeiües, such as aqueous solubility,
vapour pressure, and the octanol-water partition coefficient.
For dilute aqueous solutions within a dosed air-water system, it has been
obsenred that at equilibrium, the ratio of the aqueous and air concentrations
stays essentially constant. This ratio is defined as the Henry's constant, or in
relation to compounds containing iodine, the iodine partition coefficient (IPC).
Unfortunately, in term of determining physical properties, other than
iodomethane, organic iodides have been relatively ignored.
Hasty [4] determined the IPC of iodomethane from its concentration (c')
in a known volume of gas O/T) where no solution was present, and its
concentration in the gas phase (C) when the containei of volume VT had a
volume of liquid 013 present. The partition coefficient was defined as ...
' where CL was the concentration of methyl iodide in the liquid phase and NT and
No were- the total number .of' moh, and number of m o b in the gas phase
- m&ect@eIy.. Rearrangement of equation (2-l).gave equation (212). - --. . - . - - - - . -
A stainless steel colurnn packeâ with Chromosorb P separated methyl
iodide fmm other gases (N2, H20). For the detennination of concentrations, a
gas chromatograph was used.
The system was maintained at a spe~ifie~l temperature by circulating an
ethylene glycol-water mixture, from a temperaai~ntrolled bath, through the
jacket of the cylindrical vessel. To obtain a desired concentration of methyl
iodiâe in an experVnent (to check the concentration effed), appropriate dilutions
were made of the saturatecl methyl iodide-air rnixtwm.
Hasty reported that between 4.8 and 68.50C, the IPC for iodomethane
can be represented by equation (2-3),
where T represents the tempenature in K.
Glew and Moehnryn-Hughes [5] discovered that the IPC for iodomethane
decreased slighüy, as the air and water concentrations were increased. For their
study, the IPC was detemineci by measuring the quanüty of vapour absorbed by
a known mars of water at fixed temperature and pressure. Liks Hasty, an
equation was developed for the pradiction of IPCs, ., . . . .
where the IPC is in units of (mm Hglmolarity), and temperature is in Kelvin.
Hasty's cornparison of obsewd and calculatd values is summarized in Table
Nishkawa [6] and Postma m also measured the IPC for iodomethane. The findings of Postma are not available in this repoit. Nishizawa identifieci aie
same concentration effect previouclly stated by Giew and Moehvyn-Hughes.
Table 2.2: IPC h r idomethane C6l 01s Conc. 0.089
Tempemain m Aqmow Conc. m] 0.58
Wmn et el. [8] deteminecl the IPC of iodofm, wing the ratio of vapour
pressure and aqueous solubility (section 2.1.2.3). The vapour pressure and
aqueous solubility of iodoform wem measured in the temperature range of
293.1 5 to 353.1 5 K. The Indings a n summarired in Table 2.3.
Tibk 2.3: IPC muMing h m ratio of aqueous rolubility and vapour pnuum for iodotomi (CHI,)
L
Tempmtun (10.2 K) - Solubility (x 104 molll) 1 Vapour Pnuui. (Pa) IPC 293.15 3.92 + 0.09 1 0.166 I 0.001 576
For the experiments, vapour pressures were measured in a glass gas-
saturation apparatus. A canier gas Stream of pure dry nitrogen was saairated
with CHb vapour at the specifled temperatum. The vapour was then recovereâ
by condensation, and analysed by UV-visible spedrophotornetry, after
dissolutiori'in a k r o h volume 'of heptane. -
The solubility ineasumments were 'made with a packedcoluthn
- apparatus, consking of a glass cdumn filîeâ with solid iodofonn. Water a a
contmlkd . . . &perature was circulateci in a doseâ cimit, airough the cdumn and . . - . . *. . .
a fiow-through specaOphotornetric cell. The concentration of CHIJ at saturation
was then determined, again using W-spedrophotometry.
A later study, also co-authored by Wren [g], presented the lPCs for a
group of organic iodides. Hem, the lPCs were detemined by withdrawing
aqueous and gas samples Rom solutions within temperature controlled cells.
The analysis was performed using high-perfonnance liquid chromatography and
gas chromatography. A summary of the msults are shown in Tabk 2.4.
Taie 2.4: IPC valun marsund rt 288 K m
lPCs for iodine
, loâoacetic acid lodoethanol
Mokoukr k d l n m 03
have been measureâ diredly by Eguchi et al [IO], and
500000 71000
Toth et et [il]. Eggleton [12], followed by Pamly 1131, used measured aqueous
solubilities (14) and vapour pressures (151 to calculate IPCs for 1,. Estimated
values were also provided by Funsr and Cripps [le]. The puits of these studies
are summariied in Table 2:5.
Evans [ln using the ratio of aqueous solubility and vapour pressure reported 1.9 as the IPC foi monatomic iodine at ambient conditions.
tionation
readion to fomi hypoiodous acid (HO!).
I2+N20+ HOI+I'+H+ (2-5)
Burns et al. (181 mported HO1 to be non-volatile. In contrast to thb
theory, two studies have been perforrned with the objective of quantifying the
existence of HOI in the gas phaw. Wren and SanipIli [1Q] did observe a signal.
Nonetheless, they conduded that with the poor signal to noise ratios observed,
and the powibility that HO1 couM have ôeen created within the rnass
spectrometer, that thek data was not inconsistent with the theory of zero
volatility. Toth et al. [Il] used Whrisible spedrophotometry to try and detect
HO1 in the gas phase, but wen unsuccessfbl, and could only provide an
eatimate for the IPC based on the detedion of H06r in the liquid phase.
IPCs have .been masumd for HO1 though, through the um of vaflous
techniques, but not without a large variance in the m u b . Previous studies
Creuse p2]. Pelletier and HemphilI[23], Lin 1241, MuMer 1251, Wren and Sanimlli
[19]. a@ '6th. [Ill. Afaiough not masurad dimctly, Evms e t al. [26], along,with .
Furrer and Cripps (161 providexi esümates for the IPC of HOI. The data is
summarired in Table 2.6.
Narrel et el. [21] examinecl the conœntretion eflsct on the IPC of HO! by
rable 2.6: IPC data for HOI
exarnining concentrations vatying fhm 2.0 x 1 to 7.9 x I O 4 M.
Table 2.7: Concontmüon ehct on IPC of HO1 i t 20°C, pH 8 (21)
IWW] 1 IEm
700
Temp.[%J Ambient
I O 20 21 22
IPC[lr] iPCp3J ICCpo1
300
Expriment 1
25 30 50 70 72 75 80 1 O0 125 150 200 250 300
1
1900 I: 510
i P C w 1iPCPûJ
1 .ME4
mTm,, M 7.E%
4S8B
1 .
2.43€2
69.3 27.7 14.1 8.46
30
IPC
1090
. --
7.05E3
2.9SE5
IPCtlW
IO00 750 560
-iPCEl] 1 I P C m ~3E4
1 700
- -
--
620i320
161 k26-û
* E l 4.6€+03
7.2€+02
1.7EM2
I
5.5€+01 2.4EM1 1.3€+01
Various predicüon methds have demonstmted the existenœ of a strong
nlationship between molecular structure and the Henry's constant. A common
problem with these techniques though, is that they don't allow for a clear
interptetation of this reiatïonship. For exampie, one may argue that the use of
empirically derived contribuüons/c08nicients ignores the process of phase
distribution, and the chernical properües affecthg that process (i.e. cavity
fornation, solvent reorganization , hydrogen bonding , dipoledipole and dipole-
induced dipole interactions). M e n a method relies more on regressional
analysis than on information relating to the specific system, correction factors
may have to be addd when compounds outside of the study group are
encwntemd. It should be naed though, that with eny estimation technique,
there will always be uncertainties with respect to compounds unlike those used
in the training &S.
Up until now, outside of a b w compounds, experimentally detemined
IPCs have not been readily available for use in the development of preâicüon
methods. So a few modelling techniques were compared with experimental
msub, fOr the purpose of checking their mlative acairacy, and Mere possible,
strsngthening these models with respect to organic idides (section 4.3.3).
Aquracy, bmvity, and ease of use wem among the objectives in choosing -
madelüng techniques to examine:
A series of descriptors, originally proposed as branching indiœs by
Randic' [28] were later developed and fonnalized by Kier and Hall (29,301 under
the narne of rnolecular connedivity indices (x). The parameters are calculated
purely fr& the rnolecular structure, and are undemtood to encode
physiochemical properhies without any experimental input. Through the use of
different structural algorithms, dïfferent levd wnnectivity indices can be
calculatd (Le. Cmt, second order etc., or "ordinary" as opposed to 'vaknœ").
For example, one could calculate an ordinary zero order connectivity index ex ), along with a first order vaienœ ('$) connectivity index.
Nimalakhandan and Speece 1311 calculated connecüvity and polaritability
indices ( a) for rnany oganic compounds. The pdarizability index was defined
by equation (2-6),
@ = 4.063(# of CI) - 0.361 (# of H) - 0.767(# of double bonds) - 2.620(# of F) + 1.474(# of 1) -1.24A + 1.014K + 0.636(# of NH2) + 0.833(# of NH) - 1.6@5(# of N02) - 1 -8230 (24) . where A, K and O are indiciton (= 1 if pmmnt, = O if not) foi hydmgen bonding,
.for aldehydes. ketones, and dioxine ~ ~ e ~ y . Colledively, the parameters
were suggested to aaccount for both solute:sohrent interactions and the cavïty
formation phenornena. The parametem wem then related to Henry's consmnts
airough mgmion analysis. .ln a Iater paper, Nimelakhandan, Bmnnan and
Speece [32], attempted to make an improvement on the model, by using a larger
study group of compounds, the resulong equation ie shown below.
1-H = 1.29 + 1 .005@ - 0 . ~ 8 ' ~ " -1 .2581 (2-7)
Whem H L the dimensionleos Henry's constant ([ ]$[ lxv is the first
order valence connedhrity index, @ is the polaMbility index, and I is the
hydrogen bond indicator; I=l if aie solute participates in hydrogen bonding, !=O
othemise. The authors also attempted to take temperature into account with the
following equation.
where ...
A = 14.97 + 5.78 OX - 11 .79'Xv B = 4348 + 947 Oz - 1855'~"
Conndvity and polarkabiiii indices aie provided for a set of organic iodides in
Appendix A.
Hine and Mookerjee [331 fint introduced the method of using bond
contributions to estimate Henry's law constants. Meylan and Howard later [34]
improved on the same modelling technique. The underlying theory is that the
subunits of organic compounds have esrrentially the same effect on air-water '
equilibrium, regardless of the chernical in which they occur.
H H H H H H 1 1 1 1 d-è-l H-c-C-C-C-I
# # 1 1 1 1 H H H H For example, as shown above, it was aswmed that the quantitative contribution
of the C-l bond to air-water partitioning is the same whether it occurs in
iodoethane of 1-iodobutane. Again, using iodoethane es an example, the
following type of equation is used to calculate the IPC. * .
Log ( P h ) = 5(GH contributions) + (CC contribution) + (CI contribution) (2-9)
According to Hine and Mookerjee. this idea goes ba& to 1935, when
Butler et al. 135,381 showed that the, free energy of transfer of aliphatic '
compounds from the gas phase to the aqueous phase was approximately an
additive fundon of'tfm g r w p piesent in the wmpwnds.' To alleviate-errom - , . . - - - caused by special intemolecdar intemctirons, Hine and Moakerjee suggested
asing 'larger subgeups (ex- use --C02H, as o p p d to -CO and -OH). For
. . . moleCules . .in which there - . ais mulapie- . polar _. . subunits. moiecular interactions wn - . . . - - . . - .
. - . . . . - * * . -
get quite cornplex, and thus, the esümated Henry's constants using this method
are less accurate.
S. 1 .P.S WlmrUon band on œ q w o u r rolubllMy and vmpour pmmmuiu
One method of determining lPCs involves defining it as the ratio of
aqueous solubilii and vapour pressure. Thilr methoâ works if a few
assumptions are shown to be bue.
The first assumption states that the molar volume of the aqueous solution
is not appredably altered by the organic solute. Since aqueous solubilities of
organic iodides are very srnall, this assumption isn't really of concem.
The next assumption stipulates that the gas phase should behave ideally,
i.e. the equilibrkim partial pressun above the solution is equal to the vapour
pressure of the pure compound at the same temperature.
The third aswrnption is that because of the small solubilities being deait . *
with, the aqueous solution-activity coefficient y, doesnt change appreciably from
infinite dilution up to the solubility lima. Or in oüier muds, there is negligible
solute:solute interaction, even at saturation. '
. One schml of aiought [37'j .argua th& no maBr how dilute, due to the .. - -
bond-percolaîion nature af th.e 'hydrogen bonds wiaiin w&r, solute molecules
wiII . - YeeP each orner even et infinite distan.cea. This aawmption can be cheçked
by examining aie activity.co8fficienis as a fundion of concentration.
The most amrate way to do this, of courrie would be to look at
experimentally detennined advity 008nicients. But these values are hard to Cnd.
even for hig hly studied compounds.
To detemine acüvity coefficients, the group contribution concept UNIFAC
[38] is commonly useci. UNIFAC uses existing phase equilibriurn data to predict
equilibria for systems for whidr there s no exprimental data. This is done by
evaluating erperimntal data to obbtain paremeters for a set of functional groups,
which represent a large numkr of organic compounds. One restriction in wing
UNIFAC though, is that the appropriate software is needd beforehand.
In UNIFAC, the logarithm of the activity d c i e n t is assumed to be the
sum of a combinatorhl part y , which is a hindion of the ( 3 the molecule, and a midual part y , which is a fundion ( 3 between fundional gioups. Thus, for mdecule "i" in solution.
size and shape of
of the interactions
If the infinite dilution acüvity coenicient is known, the Henry's constant, in
tems of vapour pressure divided by aqueous concentration, can be calculated . .
- diredly with the following equation, . .
provided that you know compound's vapour pressure when airhater equilibrium
b establtshed, along with the molecular weight (M) and density (p) of the
aqueous solution.
With respect to unœrtainty, by using vapour pressure and aqueous
solubility, or any other properties to calculate the IPC, the enon associated with
these properües will be propagated in the calculated IPC. Appendix B has the
aqueous solubility, vapour pressure, and resulting IPC as a fundion of
tempemture for CH13(s) as detemined by Wren et. al. [8], Appendix C has the
same data for other various organic iodides.
The IPC is closely related to the vapour pressure and aqueous solubility
of an organic iodide. So to fully underatand the effect of temperature on the IPC,
one has to first understand its e f k t on these two promrties.
Equdion (2-1 2) (derived h m the Clausius-Clapeyron equation, see
secüon . . 2.1.3.1) can be useâ to express the tmprature dependence of vapour
pressure (for a liquid compound).
88t By analogy, the tempemture dependence of yw (= l/xmt)
- . . <
Gmpound can-6 expr68884 by equation (2-13).
for a liquid
AH =CES8 SAT - I ~ Y - SOLUnoN + constant RT
Let IPC' be defined as
IPC' = aqueous concentration - *IV =lobn - 0 œ 1 P
(2-1 4) vapour pressure Y w x p Y w vsoiution
where x is mole fraction in the liquid phase, v,, is the molar volume of the
solution (essentially water's molar volume), yw is the activii coefficient of the
compound in aqueous solution, and P is the vapour pressure of the pure organic
liquid. Assuming that the ideal gas law (IPC = (IPC') RT) applies, and that the
effect of temperature on the molar volume of the solution is negligible, the
equations describing solubility and vapour pressure as a function of temperature
can be substituteâ into the above equation to obtain an expression for the
temperature dependenœ of IPC'.
EXCESS AH Vw - AH
In (lPC) = SOLUTION R
This equation also works for gaseous .mmpounds. For solids ...
EXCESS
The enthalpy "spent" meMng canœls out, so equation (2-15) b also applicable to
EXCESS solids . The terni (AHvw - AH ) was defind by Schwatzenbach 1391 as
SOLUTION
'AH~m; the energy involved in the process of a liquid solute moving frorn the
is look at the how the enthalpies are associatd with air-water quiiibrium.
RI (vapour) t, RI (liquid) ; - A H v ~
EXCESS RI (liquid) t, RI (aqueous) ; AH- - - -
SOLUTION
EXCESS RI (vapour) o RI (aqueous) ; AH,, , AH
SOLUTION O AH~w
By subsütuting AHH,,, into equabion (2-1 5), and recognizing that In(lPC') will
have the same alope as In(lPClï) one obtains get equation (2-16).
As expected, equation (2-16) predicb that like aqueous solubility and
vapour pressure, the IPC will follow Clausius-Clapeyron behaviour. TQ solve for
the constant, one just has to use a known IPC. After lPCs were detemineci at
variow temperaturne, the resub m#s used to check if the IPC does in fact
. - follow CleusiukClapeymn type behaviour ( d o n 4.2.2). -
A theoretical relationship between the vapour pressure (P) of a pure
substance. and the temperature (T), is the Clapeyron equation
where v, and v,,~ are
iespedively, and AHvw is
the specific molar volumes of the vapour and liquid
the enthalpy of vaporization. Unless the pressure is
exûemely high, the specif~c volume of the liquid or solid is negligible relative to
that of the vapour. Therefoie it is assumed that the change in volume upon
evaporation or sublimation may be approximated by the molar volume of the gas.
- "LIQUD M SOUD - "GAS (2-1 8)
Secondly the vapour is assumed to obey the ideal gas law
(so that v, = RTIP). So, by using the above assumptions, and after
rearrangement with the aid of calculus, the Clausius-Clapeyron equation is
where R is the univemal gas constant. Below the boiling point of a substance,
AH, is a weak fundkn of temperature. 54, con si der in^ the boiling points of
ocganic iodides, and the tempkrature range king studied, AHvN is assumed to
be constant Integmüon of the ~@usius-hpeymn equation with ihe assumpüon
of a-anstant . . . AH,,, yiel(l6 eq'uath (2-20). . . .
In? = - AH VAP + oons tan t RT
To estimate the enthalpy of vaporization for the organic iodides studied,
vapour pressure data along with equation (2-20) was used (Appendix D). InP vs.
1 K was plotted, and AH, was detennined from the dope of the graph.
As with a lot of chemicals, excess enthalpies of solution for halogenated
hydrocarbons aren't generally accessible. Because of the lirnited availability of
these values, it is sometimes assumed that within a class of chemicals,
compounds of similar sizes will have similar enthalpies of solution 1391. The
valklity of this pmdœ is baseâ on the fad that one of the major components of
the free energy of solution L the cavity tem. In fa&, thb free energy is usually
considered proportional to the volume (or more accurately the su- area) of
the solute molecule. Section 2.3.2.1 provides a more indepth discussion of how
moiecular size relates to solubility.
It shouM be kept in mind, that by relating the enthalpy of solution directly
to size, unpredicted dissolution effecb, that may be spedfic to certain structures
. may be overlooked.
Aiso. by using a method of estimation based on
1401, the obsewed dependsncy~-of-entha~~i& of solutiOri
size, as mentioned in .
on the type of carbon
that won? be a significant hctor conceming organic iodides, since Mine isn't
Using experimental muits, aie estimation of exwss enthalpks of solution
is described in more detail in 88Cfimon 4,4.
Appendix E contains vapour pressures of organic iodides obtained from a
literanire sunrey. For methyl idide at 1-r temperatures one can look at the
study conducted by Wrem and Vikis [41]. Foi I, gaseous solubiliües can k found
in Tabk 2.5 (secüon 2.1.1 -1)
.The most *ammon technique useâ in the estimation of vapour pmssurss . .
is the Antoine q u a t h . The aswmption of a constant A Y , gave equatkn
InP = - +constant RT
If the temperature range is enlarged, the nt of experimental data is improved by
the introduction of a parameter C, which is used to correct for the temperature
depndenœ of AHVAP.
Equation (2-22) is the Antoine equation. The constants A, 6( AH,,dR). and C
(which is sornetimes negative) are compounddependent constants, and T is the
temperature. The Antoine equation parameters used by Stephenson and
Malancnmki [42], and Ohe (431 for varkus organic iodides are given in
Appendix F.
For otganic iodiûes for which parametem wemn't available, a method
using the boilîng point of the compound was ernployed. For this method,
Schwarzenbach et el. (391 used the Clausiur-Clapeyron equation, Kistiakowsky's
expression [39] and Çishtine4s correction factors [Ml to yield equation (2-23).
* where P is in atmosphetes, Tb is the boiling point and T is the temperature (both
in M n ) .
debyes. .
For polar derivatives of benzene, Fishtine suggested that . :
2w100, where is the dipole. moment of the derivative in uni& of
Foi naphthaleRe derivatives lie suggeaed KF = 1 + p/100.' For ' . . . . .
phenols, Fishtine suggested using KF = 1 -1 5.
. - ' In hi* report, Fiahth dedbd~mahod8 for calçulating dipole moments. - - . .
the addition of vectors lying in the same plane. Specitically, each substituent
dipole moment (p) was considerd to have a vertical (p,), and horizontal
components (vh). Each of which were calculateâ from equations (2-24) and
(2-251,
where 9 is aie angle of the substituent vedor from the horizonta il. The dipole
moment of the compound was then determineci by summing the vertical and
horizontal components of al1 the substituents. and then substituting them into
~ U & ~ O V I (2-26).
The dipok moments obtained for the organic iodides are in Appendix G. . .
By using equation (2-23) the potential for miseqtimation of low boiling .
corhpounds (i.6. s 1OOoC) is only around 5%. However, for compounds with high
. boiling points, the enor may e x d a hctor of 2.
Boiling points m m estimated for compounds for which they were - . * .
unavailabk. Vahres obtained h m ho' boiling point estimation methods weie -
compared with known boiling points, for the pu.rpose of Rnding the mort accurate
technique.
The fint method examined was the Lyden-Forman-Thodos m e W .
The methocl expanded on the extensive work done by Forman and Thodos
145,463, and Thodos 147,481 conœmi'ng the estimation of the critical temperature
&), along with other criücal properties.
Reid and S h e d [49] defined a parameter 0 = Tfi, baseâ upon the
observation that this ratio is relatively constant for many organic compounds.
A meaiod for detemining 0 has been developed by Lydersen 1501, where
various T increments relating to structural features of the molecule ware
summed, and then substituted into equation (2-28).
The critical temperature was then estimatecl with the Forman-Thodos
. method, baseci on the correlation of Tc with van der Waars constants,
where the critical tempenituie is in K, a and b are the van der ~aa l 's constants . .
- - - for the &mpound,'and R is the univemal gas constant (82.05 . -. a&3lmol K).
The constants a and b are relateâ to structure, and were also estimateci by
The second estimation method examined was Millets [50]. He combined
the work of Rackett [Hl, and that of Tyn and Calus 1521 to produce an estimation
method for the normal boiling point of any organic compound. The meaiad used
the criacal pmpeities of the compound; i.e. the critical temperature (Tc), pressure
(Pd, and volume Olt), along with 0, the ratio mentioned pmviously.
1 - - p L 2/7 (2-31 )
h-el
To solve for P, and Vc , pnwrure and volume incremenb pertaining to functional
groups wem sumrned. The surnmations wem then substituteâ into equations
(232) and (2-33).
- . whem M is the moiecular weight in gram per mole. The n o k l boiling point
. .
Appendix H compares the two meîhods, it was bund that the Lydersen-
For a solid with melting point Tm (K), one way of mlating the soliô's
vapur pressure to the subcooM liqukl's vapour pressure is through the use of
equation (2-34) given by Pmusnih [53].
AS&) is the entropy of fusion at the melting point, and R is the universal gas
constant. AS&,) can be calculated from experimentally measureâ enthalpies of
b i o n as discussed by Miller [Ml. Or, for rigkl (n~n~aliphatic) m o l ~ l e s , end
their long chain derivatives, Yalkowsky [55] suggestd an expression based on
Walden's rule to appmximate AS&,),
AS [T,) = [58.5+10.5(n-5)] (~-rnol-' -K-') (2-35) . MELT
. .
where n is the total numbew of flexing chain atoms (exclusive of protons, and
setüng n = -5 for n < 5). The equation has wide application, since 1 works for . .
many envimnriientally reievant chernicals [30]. Another meaiod of estimating
S h ) ir to masure the solubility of the rolid In a solvent in which the chernical -
ir believeâ Io form an ideal rduüon [m. The solubility of the-solid (moüL) is - . - . . . . . - . . -
then Fhs! whem F is the aga* redk ( P w ~ f 1 8 ù b ~ ~ lhu& and vs is the
sdvent m~lar volume, To mate an ideal sololution for organic iodides, bentene
1.3. PREWOUILV DCCUMINPD AQU WUS SoLumIUTIm
As done wRh vapout pressure and the IPC, a litefature survey was
performed for the aqueous solubilises
listed in Appendix 1. Solubilities for 1,
2.1.1 A )
of organic iodides. The solubiîiies are
can also be found in Table 2.5 (section
A few modelling tediniques were looked at, and comparecl with literature
values, for aie purpose of detemining the most accurate among them in relation
to organic iod#es.
organic solutes, solubility decmases with increasing sdute size. To better
understand this trend, the contribution of the separate steps in the dissolution
pmœss to the exœss enthalpy of solution can be explained, as was done by
Schwanenbach et d (391.
It should fint be understood that within an ideal solution, al/ rnolecular
interactions are identical, consqwntly, the exœss enthalpy of soluüon is zero.
First step; The organic solute molecule must be isolated hom its pure
liquid phase (without leaving a cavity). Therefore add AH1 > O, to oveicome
solute-solute attractions. For organic molecules AH1 = AHvapo~tion. Next, a
cavity that can accommodate the organic solute must be creatd in the solvent
(water). So a second enthalpy tenn, AH2 . O, must k added to overcorne the solvent-solvent attractions. The value AH2 depends of the size of the cavity (i.e.
the size of the organic solute), so this is a very important terni when dealing with
large organic solde rnokuies. 'As an aside, aqueous cavities -aren't solely
created through the introduction of a solute. Water is oRen viewed, under
normal conditions, as a fluid which contains, or has a predisposition to fomi,
Second $tep; Now the ocganic solute is ûans@rreâ into the cavity. T h
. - suréace of the organic. mokcule experien- intemokcular. attractions to the . . * .
Gmunding water molemies.. ~hese attractive forcecl indude b n def Waak and.
Înduced di* intemctions. Al&, . if the holeaik' contairis any. polar functional . . . - . . . - .
groups. dipole-dipok and ma* hydrogen-bonding efbcts will be YeW by the
solute. These atbaidive fiorces msult in a negative AHJ. In an ideal solution
- AH3 = AH1 + AH2, i.e. the enefgy gained equals the energy spent. Shinoda [58] mbrrsd to the net exœcw enthalpy change rewlang from the
introduction of the organic solute into the aqueous cavity as AHaV.
Third step (for nonideal soiutions); the water molecules immediately
surtounding the oganic solute have strong interadions with neighbouring water
molecules, but only on the side away from the organic molecule. This effedively
'solidifiesn the orientation of the shell of water moleaika sumnding the solute.
This "freezing" effed causes an enthalpy gain of AH, (AH,).
So, al1 the enthalpy contributions to the excess enthalpy of solution are
relateci to the surfaœ area of the solute. More specifically, as the nonpolar
DtcESs surface area of an organic solute incisases, so ddoes AHamN. Or in other
words, it gets harâer for the chernical to dissolve Ïn water. .
This was evident when Mackay et al. [51] correlateci aqueous solubility to
molar volume for n-alkanea and ammatics with equation (2-38)
whek S is the solubility in.moum3 andk is the molar volume in cm31mol. To use
Since the degree of contact made between the solute and solvent is
directly proportional to the solute surface area, surf&- area is intuitively superior
to molar volume for understanding solubihty. Molar volume, by definition,
encompasses portions of the moleark that are not exposed to the surface, and
thus do not participate in the 8oIute:solvent interaction.
So, ideally, maybe molecular ruhc8 area should have been used in the
correlation. Unlike molar volume however, molecular surface arsa cannot (yet)
be experimentally measured. Various studies, including Hermann [59,60,8 I], and
Pearlman [62] (who modified Hemann's work) have been involved in developing
computer programs for calculating molecular surfiace area.
Another argument in defence of molar volume was made by
Pearlman (821. He made the observation, that outside of spherical molecules
(ex. inert gases, globular proteins) which show the expected relationship,
there is a linear relationship between surface area and molar volume.
' .
The aqueous activity meMcient of an organic solute represents the many
in@racbkns . . that - takeplace between a @lute's funHonal groups and the water -
molecules sumbunding a. As 'm&tioned Won, techniques' like UNlFAC [38] . take advantage of this fact in ekrb to modet.. activity coefficients.
Consequentially, sin-a the activity coefficient . represepts. . all the .interad+
takng plaœ in a solution, they can be correlatd to aqueous solubility, as done
for example, by Arbuckle (631, and Banerjee et al. [W.
hmann [65] might have been the fimt to us8 group contributions in the
estimation of aqueous solubility. But the approach of Wakita et el. [66] is
consideml to be an improvement over Irmann's method. In Wakii's technique,
aie fragment values wem divided into lquid and solid ftagments. Regression
analysis on the derived mgmetnt values (fi) aien gave
log S = -0.957Cfi - 0.048 (2-39) for liquid solutes, and
l ~ g S = 0.063Cfi - 0.208 (2-40) for solid solutes.
The oktanol-~ater p a r a t k n ~ n t is defind by equation (2-41) '
where Co s the concentdon of the solute wMin octano1 and C, is the
concentration of the solute within water. y and v repmnt the adnrity coefficient
and molar volume mspectively.
The assurnption can be made that the activity coetficient of the
compound at dilute concentrations in soivent-sahirateci water cen be estimated
by the activity coefficient of the rame compound at saturation in pure water [3Q
This assumption itnplm # a even at saturation, the interaction of solute
molecuks with other solute mokcuîes is negligible. Since the organic iodides in
question have low aqueous solubilities, their adivity coeficients can probably be
considered independent of their concentration in the aqueous phaw.
Wth this assumption; y g w a n be mplaced by y W s a ~ ~ t ~ v W = I/c,,,,s~~
giving equation (2-41).
€$taking the logarithm of the equatky (2-42) we-an Mwith equation+(2-43). .
- M e n the solute is dilute enough, it can be assuir#d that it has no impact . - - . - . -
on. the molar volume ofthe. water-saturated octanol phase, thwr the rnolar
= 0.12 UmoL One drawback of equation volume of the octanol solvent (vodinO3 .
what this equation is often used fw, is to estimate the adivity coefiicient of a
compound in the water-saturated odanol phase ( yJ.
More di& conetations of K, to aqueous solubility wen given by
Hansch e t al (671.
For alcohols: l~g(l/S) = 1.1 13 l o g b - 0.926 For ketones: log(1lS) = 1.229 l o g b 4.720
For alkyl halides: log(1lS) = 1.221 logl
i .S.I.5 mrrdwtlon to oonnœmtîvlty Ind8oem
Nimalalakhandan and Speece [6q modelkd aqueous solubilities (molll),
with the assistance of oonnecbiv'iry (see 2.1.3.1) and polarizability (a)
indices. These paremeters wen, suggested to rspment both solute:sohrent
interadions and the cavity formation phenornena.
log s = 1.464+ 1.662'~- 1.367 O i + 1.001 0 (2-49)
As mentioned before, connectivity, along with polarizability indices for a
representative set of organic idides are provided in Appendix A.
All of what was 'dicrcussd previously conœming the estimation of
solubility applies to liquid rolutès. The solubilities for solids can be estimated in
much the same way as their vapour pressure. For example in section 2.2.2 of
this report, substitute S, for P,, and S- , f# P--,, where . S is aqueous solubility.
Appendix J shows aie performance of each of the above methds,
in compaflson to litemhre values of aqueous solubilities. Hansch et el.% [67]
correlations of & to aqueous solubility were deteminecl to be the most accurate
in relation organic iodides. This method was u s d to calculate aqueous
solubilities for use ( w h vapour pressure) ln predictkig lPCs (Appendix C). Of
course, in looking at the uncertainty, aie emr inherent in the calculateâ
solubiliies (and vapour pressures) was propagated in the estimated IPCs.
The most widely us& methods for detennining Henry's constants
ewperimentally are equilibrium pailitioning in closed systems (EPICS), which is
based on a stationary approach, and batch air stripping (BAS), which is baseâ
on a dynamic approach developed by Macûay et al. [3].
In EPICS, equal arnounts of aie organic compound are placed into two
diffsrent liquid volumes, msuiüng in two difbrent concentrations. The two
solutions are equilibrateû with air, and the Henry's constant is üien detemined.
h m the headspaœ concentrations, and a subsequent maso balance on the h o
systems. There are some systematic problems associated with EPICS though,
including a decreased sensitiiity for lower Henry's constants, adsorption of the
- compound ont0 large intediacial amas, and absorption (by diffusion) into
. . For this thesis, .a variation on the BAS procedure was developed (the
* . - . . - procedure is discussed in more detail in section 3.3). The original procedure
. . d e v i d . - . by Mackay and his a d a t e s involved bubbling ai? thrdgh an aqueous - .
- solution, and then aedting the concen6ation of the organic solute, with the help . -
* . of UV s ~ o h w n r t r y . A proposed maaiematical description of W stripping . -
where C and Co are the liquid concentration and initial liquid concentration
n#rped*veiy, H is the Henry's constant (atm m3/mol), G is the gas flowrate (m3/h),
V is the vesse1 volume (m3), R is aie gas constant (m3 atm g -mol-' K'), T is the
system temperature in Kelvin, and t is the time (II). Using this equation, the
loganthm of the concentration was plotted against time, and the dope was used
to determine the Henry's constant.
One of the setbacks of this piocedure is thet the organic solute of interest
had to be below its saturation limit, or else the compound would just keep
dissolving as the solutkn was stripped.
The new procedure used in mis project involved the same stfipping
process, but the novel aspect hem was the use of 1-131 labelled organic iodides.
Concentrations .for. each phase at equilibrium, in @mis of acüvity per volume
were measured, and then from the ratio of these values, lPCs were detemined. .
The use of thh radiochernical app&ach not only enabled both phases to' be . .
'analysed di~8Cf(y at saturation, but it al80 allowed lPCs to be measured relatively
Considemg the adivities and -total volumes k ing deait with. radiolysis
was mot a concem. Also, because of the manak useâ within the stripping
column. abeorpgon could be acuurrned to be negligible, and due b the tuibdent
conditions p~@ent, adsorpüon was unlikelîy. - -. . -
The fi mt step of the expeiimental procedure involved radioadively tracing
the organic iodide via atom exchange with the radioadive isotope "'1 (in the
fomi of Nal).
This was accomplishd by mWng 200 PL of the organic iodide with 50-100 PL
(50-150 pCi) of tracer. The acaial amount of tracer useâ depended on both the
aclivity present, and the extent of atorn exchange exp8cted. Acetone (lml) was
used as the solvent for this reactrton, because of Q ability to dissolve both the
polar sodium iodide and the (wually) non-polar organic iodide. The extent of
atom exchange varied according to the organic k i ng traced, .in general the .
organic iodidets were observecl to have activities in the range of 5-50 pCi.
One day was allowed for suffident atom exchange to take place. After .. .
this, 250 ml of a 1:4 by mass solution of NaCI was addeâ to the system to - .
effectively 'sa& out" the organic iodkle. The'dution was cehtiMigd, and the . .
predpitated organic iodMi wao then ~kpafated, and p l a d ha water. l'hi8 new-
seilution &en shaîcen, and &nWbged b separate'any r&iaining NaCl
water. The amount of impufities pmsent in the saRed out phase was negligibk
Henœ, the washing phase was a precautionary, as opposed to necessary stage.
AftM the tmced organic iodide was placed in the stripping vessel, the
vesml was shaken to help ensum a homogeneous solution.
Sinœ the sodium iodide tracer, and 24odoethanol were miscible, the two
chemicals were mixed directly (sans aœtone), and left for a day to allow time for
atom exchange. To üien separate the 2-iodoethanol from aie tracer, the mixture
was washed with chlorofom. The chlorofom phase was aien separateci and
washed with distilled water. The distilîed water, now containing labelled 2-
iodcmthanol, was then plaœd in the stripping column.
For molecular iodine, the tfacing todc place right in the stripping column.
The cdumn was fmt filled with a Murated ( rn I O J M) solution of 1 , -and tracer
- was th* added di- to ' the cdumn. ThW. mahod kqt the aqueow 1, . . .
conœntrdon et a maximum, 'to avoid thk formation of HOI. A .W-visible -
specbophotometer wacr used to track the iodine con.œntratîon within'the cdumn.
Uring the following madon scheme, sodium hypochlorite, along with
tracer was used to prepare labelid HOI.
NaOCl + ID + Na* + 01' + CI' (3-3) 01- + H+ (frorn aqueous soivent) + HO1 (34)
Tham fint Wo mactions am efliectively irmtantaneous. The concentration useâ
for the limiting magent (I') was approximately I O 4 M. If aie concentration was
allowed to go over IO*' M, other reaaions, like the follawing would have becorne
a conœrn.
HOI + HO1 + HIO, + 1' + H' ( k s 22 moV(L s), [69,70] ) HO1 + H l 4 + 10i + 1' + 2HI ( k o 240 moV(Lm s), [71] )
Sma 0A"H UR 8TRiPPiNa PRW-
A schematic of the aimow within the apparatus is shown in Figure 3.1.
An aspirator was used to mate aiflow through the system. Air was first bubbkd
. through a vesse1 containing on@ water. By' saturating the air before it reached
the main stripping vessd, evaporaüon of water from the arganic idide solution
was minimized. This was .important, since when a balance on the total activity
w ~ s perfomed, it was assumed the the total volume of the solution stayed
constant
After. the air was humidifkd, .it was bubbieâ through &e organk iodide. . .
solution.'The bubbler umd demaseci the size of the bubbles, henceh&ea&g
aie surface ama to volum ratio. This in effiect minimized mas8 transbr . . - . . - . .
limihiions, which helped to e ~ u m equilibrium- betwaen aie air and the. solution. . . . . - - . .
Packing was placed in the stripping colurnn to increase the reridence time of air
within the column, again helping the organic iodide mach equilibriurn between
the two phases.
The air was then p a s d through an activateci catbon filter, ont0 which the
volatile organic idide was adsorbed. It was assumeâ that the only volatile
radioactive species within the system was the organic iodide. This was a
masonabk assumption, since the oniy other potentially radioactive species
1311 - within the system was (non-volatile) .
To determine lPCs at non-ambient tempemtums, a temperature control
loop was utilùd. The loop consisted of a CobPamwrr Digisense digital
them~couple aiemorneter, an Omga CN76000 PID controller, and heating
tape wrapped around the cdumn. Glass wool was uwd for insulation ta minimise
heat loss.
Conceming the potential for misting, ie. moisture being canied via air(low
to the adivated carbon filter, it was not seen to be a problem. If it was,prevelant,
the range in gas activity observecl woukl not have been present. .The.
experimentai design also ensured that aie two hindarnental assurnptions
concerning batch air stripping wre fuifilied.
. ~he fint aqwmpüun states that.equilibrium b e k n the gas and .aqueous
phase is reached by the tim6'air tfavelr to the tbp of the stripping column. The . .
two main factors affecting equilibrium wiaiin a stripping cdumn is 1) the ratio (R)- . . . . . . . . .
* - of-the mqss knsfei area (Le.-surface ama) to volume fw each bubblqof air, and
the müo was kept at a maximum, and the gas flow rate was kept at a minimum.
To check if equilibrium was attained within aie system, the WG radio was varied,
by incmasing the gas f l m t e from 20 to 50 rnVmin, effectively Iowenng the
reaidenœ time of air within the column. Sinœ the sire of the bubbles did not
change with the increase, the WG ratio was effectively lowered. The adivity per
volume in the gas phase remained constant, even with the variation of mass
ûansfer variables within the system. So, it was concluded that equilibrium was
indeed attained within the column.
The second aswmption conœming BAS is that the liquid phase being
studied is homogeneous. The main conœm hem is with respect to the organic
iodides with very low aqueous solubiliies. It is with these cornpounds that
conœntraüon gradients are more Iikely to orbe wiaiin the stripping column.
For a non-homogeneous solvent, it is highly unlikeiy that mass transfer
variables .could be changed as they mm, without any notiœable change in the
pattitioning behaviour of the compounds. .
To help maintsin the homogeneity of the solvent, the column was shaken . -
peiiodically. Also, when sampling the liquid phase with a syringe, aie synnge
. - wes rinseâ a few thes by fmt m v i n g , end then injecting soivent badc into the .
cdumn. This created turbulenœ within the system, again, hgping to pmnre . .
homogeneity.
A variation of the stripping apparatus develom by Mackay et el. (31 was
u d . A picture of the apparatus is s h m in Fig. 1.
Fig. 1 : Exp.ilmental Apparatus
The stripping vessels had a diameter of 4.5 cm and a heigM of 25 cm. To allow
airflow betweem the dimient units (i.e. the two vessels and the flometer) 0.25"
Tygon tubing was useci.
After a steady temprature was reached, each expriment was fun for
approximately 20 minutes, within this tim, four runs, each lasting approximately
five minutes were perfomed. This led to four deteminations of the IPC per
expriment By multiplying the time of the run by the gas flowrate, the total gas
volume was detemined. The f l m t e was mearumcf with Cole-Panner N112-02
glass flowtubes. Becaucw, of the amounts being deait with, it was assumed that
the volume of the organic iodide vapour was negligible in relation to the total
volume.
10 masure the gas phase conœntration, an adivated carôon filter was
UA to adsoib any volatile organic iodide mokuks. To check if all volatile
molecuks were k i n g adsarbsd by the filter, two Ilteri, were pleced in series. In
doing athis., it was consistently found that the second filter had negligible ectivity
adsorbai. The adhrity of the adivatd carbon filter was detemineci through the '
- use df a Wallac 1282 CompuGamma amm ma Counte~ The activity was then divideci by the gas volume to deB.mine-a cpnœnûaüon jn te8ms of activity/ml..To . . . - . - .
-minimise eimr, the counting time was' mqximbd;' the .emr assoc ia with
counting nngetd from 0.1 % to 4%. - -
For the aqueous concentrations, Iml lquid samples were taken with R D
disposable 1 cc syringes through a sealed sampüng port at the bottom of aie
cdumn. The samples were taken at the beginning, middle, and end of each nin.
and then counted wiai the same gamma counter mentioned above. During each
an, the d m a w in liquid acüvii was found to be negligible, so the middle
sample was chosen to represent the aqueous concentration. Considering an
average value, as opposeâ to using the middle sample was shown b have an
insignifiant effect on the IPC.
To determine the fraction of adhrity due to the organic iodide, another
1 ml sample was taken. The sample was then extraded three times with
chlomfom. This fadlitatd the parüüoning of the organic iodide into the
chloroform phase,. sepamting it- h m any acüvity associated. with inorganic .
spedes. The total chkmfom adivity was compareci to the aqueous acüvity, for
- the purpose of calculating the 'organic fractionn. The adhrity measured fmm the .
(unrinsed) primary sample was~multiplied by this fraction, and divided by the
.saempk volume (1 ml) to give the aqmi" phase ' concentration. This
concantmtion.was dkkkd. by .the gas phar;e ¢ration to give the IPC. To
more ctearfy deinonstrate the cafculation meoiod, iodomethark is u.seci as an- . .
a: samples were 1 ml of aquoow solution b: acüvity in unlb of wunb per minute c: (chlorofomr adhrity)/(toW aqueous activity) d: b x c, sinœ unQ wem (wunts pew minute)lml a: (chamal Citer activity)l(volume of gar),
whem (volume of gas) = (gas flowrate) x (time of nin) f: d/e
Dunng the study of each compound, a gas chromatograph (HP 5890 11)
was used in conjuncüon with a mass selective detector (HP 5971A) to check if
any change in speciaon had 'taken place m i n the .stripping column via
hydrolysis. For each compound, a s p t m r n was created for a'sample taken . .
from the stripping column, this spednirn was then compared to one
corresponding to the pure organic iodide. For spectra that weren't availabk, .
they were createâ using pum samples, and are shown in Appendix K.
The liquid phase anaîysis was pmf6mJed by extrading the or& iodide - .. -
from the aqueouo phase into i d n e , the isooctane was then injectedBcfedinto .. the
The gas phase analyris was perfonned via wlid phare microextraction
(SPME). The use of SPME invohmd the placement an adsorbent above the
solution king studied (within a dored system), to adsorb any volatile species.
To help identify the compounds, a GCIMS was again used.
For the compounds looked at, it was found that outside of iodoaromatics
(section ?), no change in speciation took place. Wfih respect to HOI, as
mentioned before, inorganic iodine chemistry can get quite cornplex, and as such
was considerd ou ide the scope of this project. Consequently, the main
concem with respect to these compounds was to determine the ueffsdve" IPC.
3.7 CHWIWCALS USED
The chernicals used, and their suppliers are shown in Appendix L.
10 check the accuracy of the expe@nental method, experimentally
determined lPCs wen compared with thow fbund in lbrature [QI. It should be
noted aiough that literahire values can't ever be considered 100% accurate.
Aîthough there werent many previously reported values for experimentally
detennined IPCs, those that wem available were found ta be in general
agreement with those detemined using the stripping column method. Results
for al/ the wmpounds studieâ are provided in Appndk M.
Table 4.1 : Cornparbon of Measund and Libmtum IPCs Compound 1 @cporimontil (WC) 1 Ubmtun (O] (2VC)
b: value. for 2-lodophenol . .
' lodomeaiene lodoathane
. 1 -lodobutane . " 2-lodoethanol
.4-lodophenol J
... . . For nonambient condiaons, Figure 4.1 compares the lPCs mcwrured for .
iodomethane and 1, with thme previousiy. d@ennined by Glew and Moeîwyn-
. . Hughes [a, and P~mly [13]. ri3speCarely. : The kft and righfwis pertein . to - .12, and
a: value for 2-lodobutane
4.2 a0.4 3.1 I 0.3 2.7 i 0.4
3.2€+04 I 2E+03 1.3€+03 i 3E+O1 "
- 4.8 2.8
I
2.28 7.1 E+04 7.3h02
0 - O. 0- - - 1- m.: - . O - . %: : - - y O *. O* -. - . - 0 0 . * * . . - F:. - - - . . . - . _ . . . . . . - * . . . . . . . -
lnherent within the experîmntal ~ ~ ~ ~ e d u m , each experiment invotved the
performance of four nins. Conquently, it was found that the modified BAS
method consistently produced accu rate resub.
If not othennrise stated, al1 the m u b d i s c u d below were detennined at
S.T.P. Non-ambient temperotures are examined in section 4.2.2 (Effect of
Temperature). Alsol al1 mgressions were performed with a standard spreadsheet
Generally within a group of organic comimunds, both aqueow solubility .
and. vapour pressure demase. with an increak in m w l a r ske. WSa, re&ct - . .
speCmcalîy, the surface area) of the (Iiquid) solute increases, its aqwour
. - solubilii decreases due. to- the i nc rsm. enthalpy cost Co-ing vapour :
pressure, the decrease with size is principally due to the stronger van der Waals
attradCons associatecl with larger mokcules.
W~ the decrease in both aqueous solubility and vapour pressure, the
IPC might not be expected to exhibit a significant dependenœ on molar volume
(Le. molecular sue). For organic iodides though, dong with the other organic
compounds looked at (Figures 4.2, 4.3), the IPC (andlor Henry's constant)
decrscisd as aie molar volume incmad. The only anomaly to this trend was
2-bromoethylbenzene, alaiough it should be noted though that the Henry's
constant for this compound was estimated [12]. The trend suggests that for
organic compounds within a -hase system, the impact of molecular size on
aqueous solubility outweighs its e W on vapour pressure.
In companng iodoalkanes and iodoaromatics, the IPC for
iodoaromatics experiences a steeper negathre slope with respect to molar
volume. It m m 8 that the IPC for iodoaromatics is a stronger function of
molecular size.
Figures 4.2 and 4.3 suggeqt thkt the addition of an:iodine, or any
other halogen substituent, wiH effectiyely increase the Henry's constant. In
cornparing- the pioprties of saturateci alkanes to (mono-)halogenatecl alkanes,
the habgen substituent mduœs.the vapour pressure, by increasing the size of
the- molecble. ~ i a i i n the aqueout phase though, the reduœd - finity'
componding to the increase in ske ir oubmighed by the eflea of the dipole -- . .
moment cieabd by t(M substituent, uiümateiy -mulang in a higher.solubility. The . .
substituent's enect on the Henry's constant slightly incrsases with the site of the
halogen (F
The lPCs for both 2-kdoethanol and Ciodophenol were found to be high
in cornparison with aie iodoalkanes and iodoaromaücs, at approximately 32000
and 1300 respectively. This can be attributed to the fact that both of these
compounds contain the hydroxyl fundianal group. This substituent strongly
interacts with water via the formation of hydiogen bonds and large orientational
forces, resuîüng in an increased affinity for the aqueous phase.
For 4-iodophenol, the potentiel impact of pH on the IPC was also
cansided. Clodophenol, like any other phenol can exist in the 'phenol" or
'enoln type foms shown in Figure 4.4.
Figun 4.a A) phmol am, of bkdaphanol, 8) mol knn of Clodophenol . . - r .. . - .
The acid.dis&ciation constant (KJ for Cldophen01 is deined by the followirtg
equation. . ,. - . . . . . . . -
Z - [enol f o n ] ~ ' ] (Ka)4-iodopbnol [phenol forni] (4-1)
When the pH of the system is equal to the pK,, the concentrations of the
enol and phenol forms are equal. Shce the K, for 4-lodophenol, or any other
phenol iodide wasn't available, the K, for 4-lodophenol was assumed to be
similar to the K, for echlorophenol. This assumption shouldn't be too far off,
since the acidic behaviour of any phenol is primarily contrdleû by the hydroxyl
substituent. The pK, for 4-drlorophenol is 0.2, and the pH of the deionized water
used was approximately 6, hence the equilibrium for 4-iodophenol was shifted to
the uacidic ride". Consequently, 1 was the IPC of the phenol fonn that was
measured.
To examine the enol fom, soâium hydroxide was added to attain a pH of
approximately 13. The pH was measursd wing an externat pH probe (Orion,
mode1 410A). It was found that because of its charged nature, compared to
approximately 1300 for the phenol forni, the en01 form had a much higher IPC of
approximately 2 1 000. . . '
Mine (1 J was detennined to be rslatively volatile, with an IPC of
to 'be îess volatik, with an IPC of approxhnateiy 22000. Of course, with the-kw . . * L
concentrations ( < IO-' TM ) n d d to study HOI. aqueous impurjtjes pmsent a
possible source of error. During the preparation of HOI, lodine could have
potentially reacted wiai organic impunties to produce volatile organic iodides. So
theoretically, it is possible, that wnceming HOI, part of the measured acüvity in
the gas phase could have been due to volatile impuriaes. Consequently, if any
other volatile organic iodides were present in the system. the actual IPC for HO1
might be higher than 22000. Or, HO1 migM even be effectively non-volatile, as
suggested by Burns et al. (181.
As previowrly done by Wren and Sanipelli (191, and Toth et al.[ll], an
attempt was made to detect HO1 in the gas phase using solid phase
microextraction (SPME) and a GCMS. An appropriate adsorbent Zker was
chosen, and placed above a sample of an HO1 solution to adsoib any volatile
species. No signal was det8ct8d though. Assuming that HO1 is indeed volatile,
one possible explanation awld be that the dilute nature of the system
( ~ 1 x 1 Od M) didn't allow for easy detedion.
For every ,wmpound studied, as the ûarnpemtu(eq.was increased, the .. . . .. .
value of the IPC degeased. in examining. aqueow ' rdubility and vapour
pressure separately, both pcoperbes incmase with temperature. But within an
. air-water BysWm,. it seems that as the l#nperature goes up. an.organic'iodide's
the fact that e higher temperature repmnts a highr kvel of energy fw the two-
phase systern. This exbg energy enables more dute molecules within the
aqueous phase to evapomte into the gas phase. Figure 4.5 shows how the air-
water equilibrium changes for a reprewntative set of organic iodides with respect
- -
A 2-iodoethanol
HO1
eiodine
e4-iodophenol
@lodofom
iodomethane
A 1 4odobutane
, Figum 4.5: Behaviour of IPC with respect to tempmtum
In Figure 4.5, unless otherwise reported by an enor bar, the size of the
data labeb repmsent the uncertainty pmmt within the metasurement.
To review, in estimating the IPC as a hindion of temperature, it was
assumed that the IPC follows Clausius-Clapeyron type behaviour as shown in
equation (2-1 6).
Where AHHaRy npresents the enthalpy involveâ in the movement of a solute
fMn the gas phase to the aqueous phase. The proce88 is generally exothermic.
The Clausius-Clapeyron relationship was tested against experimentally
detennined IPCs, by plotang In(lPCîT) vs. 1TT for a set of organic iodides
(Figure 4.6). Appendix N provides the same test for the compounds studied, but
not represented in Figure 4.6.
A linear nlationship ôetween In(lPCIT) and 1fï was obtained, showing
that for the temperature range studied, the IPC does follow Clausius-Clapeyron
type behaviour, and the assumption of a constant AH,, was accurate in
relation to organic iodides.
Since the dope from equation (2-16) is - AHHENRy/R, and the slopes h m Figure 4.6 are positive, AH,,,rnust in fad be negative, and thus, represents an
exothennic pmcess, as expected. Frorn the experimental results, values for
AH,, were detemineci. Recalling that enthalpy is a state function, excess
enthalpies of solution were then determined by adding AH,, values to
enthalpies of vaporization [5].
Tabk 4.2: AH-lbi organic kdY#
- a: ûetecmined.thiough adâiüon d AH- a n d ' ~ b . . b: Ljtereitun value = 3.4E+04 [8]
c: Litehns value = 2.3€+04 [731
lodoethane 1 -1odobutane 1 -lodoheplane 1 -1odoodane 24odoethanol 4-lodophenol
lodofom (CHI,)
12 *
HOI' .
Compound 1
lodomethane A ~ P [Jlmoa
2.8€+04 AH- [Jlmol]
-2.5€+04 -2.1 E+04 -1 SE+M -3.3€+03 -1.5€+04 -2.4€+04 4.9€+04 -2*7E+û4
. *-2.QE+M - -2.3E+M .
AH-& ' [Jlmou 2.4E+03
3.2E+ô4 3.9€+04 5.4€+04 5.7€+04 '4.9E+ô4 9.8E+ô4 - 7.0€+04 . 6.5€+04
- Not availabk '
1.1€+04 .. 2.4€+04 5.0€+04 4.2E+ô4 2.6EW 7.9€+04 . 4.3E+Wb 3.6€+MC
For the iodoalkanes studkd, with the only exception being 1-iodoheptane,
the process of moving from the gas phase to the aqueous phase became less
exothennic as the compound got bigger. Alsol it smms that as the compound
grows in site, AH= increases bster than AHvw does. Or in other words, the
solution process gets more endothermic faster than the vaporization process
does.
As obsewed by Robaugh If4], the C--1 bond whin iodobenzene and
24odotoluene is vulnerable to thermal decomposition. With respect to
iodobenzene, Wren et el. [QI reported a hydrolysis rate constant of
4.6 x I O J (s-') at 70°C, which conesponds to a haK life of approximately two and
half minutes.
Because of aiis phenornena conceming iodoaromatics, for temperatures
above 25°C. definitive IPCs could not k obtained for iodobenzene.
2-iodotoluene. or 1-iodonaphthalene using the prescribed experimental method.
. . 75OC was tracked with-the help .of a GCMS in Figures 4.7, and 4.8. .
nc: OMINROOM.~ Abundance
Flgum 4.7: lodoknzene prk; Tempntun 76%, Time: O houn
For the sample of iodobenzene anal- at 75OC. after an hour the
spectrum abundance dropped from an initial value of approximately 3E+07 to a
value of 1.2€+07. Other evidence of thermal decomposition concerning
iodoarornatics included a decrease in the organic fraction of iodine activity. The
difficulty involved in determining iodoaromatic lPCs didn't lie in the formation of
inorganic species. The chloroform separation described in W o n 3.6.2 allowed
fw the separate analysis of inorganic, and organic species in the aqueous
phase. The problem arose in the analysis of the gas phase, the activity per
volume in the gas phase showed an appreciable variance, limiting the ability to
determine conclusive IPCs.
- Recent studies have s h m iodine to have a significant role in
stratospheric ozone depletion. One possible pathwayefor iodine, invokes the
emission of volatile organic iodides from the ocean. The compound chosen
for analysis in relation to ozone depletion was iodomethane, since this is
most likely the most abundant organic iodide within the ocean. The
conceritratkn of sodium chloride in the w a n is . . apppximately 0.5 M. .
Accordingly, the experimen-ial plibC8dure concerning the effect of salinity
inyolved lowering the concentration of sdium &bride from this amount, until - . .
no signqcant effkt was seen on . . thq IPC. Of cou , there *are other .
sulphate, magnesium, cadmium, and potassium. But these chernicals are
di lm when compared to the amounts of sodium and chloride present in the
ocean. Schwarzenbach [3Q] nported that the eror introduced when using
NACl instead actual salt water for prediding the effect of salinity on solubili
and adivity coefficients of organic compounds in seawater is only about 10%.
To se8 b Med on the IPC, the concentration of borîc acid was al80
lowered until its efbt on the IPC becarne negligible. Boric acid is used in
light water reacton as a neutron absorber.
Theoretically, via saltingout phenornena, the addition of NaCl
andlor H,BO, should lower the value of the IPC. Franks and Evans [75]
described the saitingout efbct as nonpolar solde molecules being 'squeezeci
out" of sdution by the prefeisntial atbadion of polar water molecules ta charged
ions. One can imagine the addition of an electrolyte as causing a decrease in
the =degree of crystallinity" of water. EffBCtjvely, the added ions break the
structure of water, or altematively,.decrease the mean ske of aqueous cavities.
This wnsequently lowers the capacity of the aqueous solvent to accommodate
non-polar sdutes. The cause of the structure breaking is presumbly 'the
o&ting influence of aie ionic field overriding the 'normal" structural arienting
influenm of neighbourniq wateï moleculm. . - . . .. . . . .
The observed e d b d of addlng an ebtp)yfs wasn't too , dramatic .
.(Figure 4.9). rnaybe in part due to the fact that the iodine substituent is elecfron - - . .
. withdrawing, thus creating a dipole within the mo
with a dipok moment pmsent, the "volatilizingw effect of adding ions to the
solvent is reduced.
It should be noteci though, that ignoring polar substituents, iodomethane is
the most pokr of the organic iodides, and the e f k t of salinity on aie IPC of
other 'I~ss polar" organic iodides is expeded to be more pronounced.
Sodium chbride seemed to have a larger effect on the IPC than boric acid
did. One explanation could be the fad that sodium chloride completely
dissociates, whik boric acid docw not. Thur, in comparing equal concentrations,
a sodium chloride solution is more polar than a borie aciâ solution.
For the organic iodides for wMch rneasured aqueous solubilities and
vapour pressures were available, Table 4.3 compares the IPC detennined by the
dimenrionlesr ratio of aqueous solubility and vapour pressure, and the adual
rneasured IPC.
Table4.3: Eumination of Ratio Auumption for IPC
COMPOUND [AQUIOW]~~ [Ai& MO 1 - IPC lodomethane 0.7E-02 I761 ' .. '2.2€-02[41] 4.4 4.2 I 0.4 , -
. lodoethane 2.6E-02 1761 7.3E-03 [41] ' ' 3.5 3.1 10.3
Not shown in Table 4;3 is the enor associated wiai the ratio values. which
.. belf is a fundon of the emn pment 'in the - solubility and vapour (ne& . .
The auumption of the IPC equalling the ratio of aqwous solubility and
vapour pressure reems to be accurate with respect to organic iodides.
Nonethekss, considering different substituents and their specilic properüee,
caution should be taken when applying the assumption to other organic iodides.
Also, the assumpüon can't be really applied to compounds like 2-iodoethanol,
and 4-iodophenol, rince they are totally miscible with water (infinite solubility).
Calculated ratios of aqueous solubility and vapour pressure are shown in
Appendix C.
4S.f S i e L l O N Of MOST ACCURATE IPC MODEL
The most accurate IPC estimation technique was found to be the one
proposeâ by Meylan and Howard [W], who impnwed upon the findings of Hine
and: Mookerjee [33]. lheir aieory 2.1 -2.2) that chmical bonds have the
same quantitative e f k t on the Henry's constant was testeâ with respect to alkyi .
idides. For the elkyl idides shawn in Fgum 4.10, aie only diflerence in -
molecular sttlldure lies in the number of CH, groups prerent. In the gr& kiow, . . .
this difference is iepmentd by .the 'numkr of* caMn a t h s p-nt in the . .
compwnd.. The. lirnrar-relation i n Fgure 4.10 implies that the .CH, group had 8
relatively constant etlect on the IPC (or more spdkally the loganthm of the
O 2 4 6 8 10
Number of Carbons
Figure 4.10: Examination of conabnt bond contribution ruumption - for rlkyi iodides
Akng with demonstrating the valiâity of Hine and Mookerjeets theory with
respect to alkyi iodidesI ftom this plot, an .quation can be dev,eloped to give an .
estimate for the IPC of an akyl iadide, with meiely a knowledge of the number of
This equation, of cou^. ignores the (maybe negligible) efiiect of branching in
The perfbmance of Meylan and Howard's bond contribution mode1 with
respect to a sample set of organic iodides is shown in Table 4.4. A standard
deviation of 0.34 log IPC units was repofled for the model.
Table 4.4: Comprrkon at 2 5 ' ~ of M.raumâ IPCs, and th- alculatd
Exœpt for 1-iodonaphthalene, and 4-iodophenol, the measureâ lPCs
usiG the bond contribution mothod of Meylan and Howarâ [34]
seemed to follow the behaviour prsdictd by Hine and Mooke-rJee (331 (et laast
within an order of magnifude); In. an effort to obtain an explanation for the
anomalies, the comsponding organic chlorides vkre looked at, but only
MEASURED 4.2 î 0.4 3.1 i 0.3 1.6 k 0.1
0.57 * 0.1 0.62 i 0.03
2 2 + i 1Q*1 1611
3.2E+û4 * 1700 1.3E+03 I 35 4.3€+02 I 25
COMPOUND ladomethane
L
lodoethane 1 -lodobutane 1 -lodoheptane 1-ladooctane lodobenzene
I
2-lodotoluene 1 -lodonaphthalene
2-todoethanol . Cloâophenol .
lodofom
estimated Henry'* constant8 w m avaikbk. Sinœ naphthalene, and phenol
ESTIMATED 4.4 3.4 1 -9
0.81 0.61 20 18
2.0€+02 9.2€+04 1.9€+05 8.0€+02
derivatives were used in .the .original training set, s-c interactions, inherent to
the iodine subsbihjent am proôably the reason behind .the poor acarracy -
concerning aiese two compounds.
The next s p was the improvernent of the CHI contribution developed
by Meylan and Howard 1341, using acquired experimental data. A C--1
contribution was al- derived to help account for the characteristic interactions
within aromatic iodides.
Tabla 4.5: ûetwmbtion of impmved C-4 bond contribution ) COMPOUW 1 Cm,4 bond contribution 1
As shown in TaMe 4.5, C-l contributions were detemiined for a set of
alkyl iodides. This was done by s u b t r ~ ~ n ~ al1 of the stnictural contributions
from the IPC value until only the C-l contribution remained. This method
assumed that aie contributions pertaining to the hydrocarbon portion of the . "
moleoules were acwrate, which is masoilable sinœ they wre previously tested
I
lodomethane lodoethane
1 -1odobutane L
1 -lodoheptane 1 -1odooctane
Average
against a large ~t of organic compounds (see [Ml). The contributions were then
0.9802 0,9740 0.9308 0,8542 1,0130 0.96ûS
averaged., resulting in a new, more repreaentative value for the C--1 bond . .
,contribution. WiPh respect to the sample .set of organic iodides, the previous
contribution of l.0074 gave an accuncy of I 0.06 log(lPC) units, while the new. . . r
value of 0;9505 gbm an accumcy -of 10.05 m(lPC) units. The impmverrmnt
contribution was the uncerteinty with respectto how it was developed. Thus, like
the alkyl iodides, expecimntal data was used to derive a contribUh*on. Only this
the, aromatic iodide values were usd.
f rbk 4.6: btmninrtion of improv COMPOOND 1 C -- 4 bond contribution
Because of the inaccuracy of the mode1 in relation to 1-iodonaphthalene,
it b suggested that caution be taken when predicting the lPCs of polyaromatic
iodides. In relation to monoaromtic iodides though, a C-4 contribution of
0.51 31 has been determineci, using the modest sampk set of iodobenzene, and
2-iodotoluene.
In analysing the .method of Nimalakhandan et al. 1321 (8ection -.
it was found not to be suffidently accurate with respect to organic
mis was expecteci though, rince the mdei was onîy tested within the
The next technique examineâ to preâid the IPC as a fundion of
As shown in d o n 4.22, the IPC dom follow Clausius-Clapeyron type
ln (y) = - [ A:m]' T + Schwanenbach propod using estimated (where needed)
enaialpies of soknion and the enthalpies of vaporization to detemine values for
AHmw The calculated AH,, couid then be pluggd into equation (2-le), for
the purpose of detemining aie Henry's constant as a funcüon of temperature.
The enaialpy of vaporization can be estimateci, as describecl in
seclion 2.1.3.1. Sinœ the enthalpies of vaporization were calculated primarily
using the Antoine equation, they can be assumed to be mamably accurate.
The exeess enthalpies of solution on the other hand, are traditionally predided
on the basis of the molecular sue.
The difficulty with this technique is that excess enthalpies of
solution for. lquid solutes are harâ to find, much less, thow that can be
reasonably compared to organic iodides. In searchhg br measumd enthalpies,
no compounds wem founâ that xould be reasonably compared to' organic . . . - . .
iodides. Of coune, the m u b of this project wilb be the fimt $tep in redifyng that
situation. -. . . . - .
- + Abmmon tfs] pmpmd that b r rom organic compounds, the
CH, group has a constent additive pmperty in mlatiqn to excerrs enthalpies of . .
- solutiok The saine t . kit of phenomena was obseived. with kpect to - cilkyl *
Table 4.7: Mwminatkn of quantitative CH, contribution to AH UCESS S O t m
The detemined CH, contribuüons were averaged, and a value of
- -. -
7.9€+03 Jlmol is suggested to describe the quantitative dbct of the CH, on the
Compound 1 [Jlmoil [domethane 2.4€+03
excess enthalpy of solution.
CHz contribution [Jhoîl
UCESS Also, using deteminecl AHwLmM values for organic iodides, dong
with those found for fluoromethane, chloromeaiane, and bromomethane [5], a
correlation for alicyl halides, betweten the excess enthalpy of solution and molar
volume has been developeû.
Using regmssion analysis, the fdlowing equaüon was developed,
AHL-- = 337 (molar volume) - 1.99E94 (4-3) where the enthalpy and molar volume have units of Jfmol and mUmol
respectively . UCESS The uîümate test was to see whether piedicteci AHsoLmoN and
A b values could be u d to mmod the IPC as a function of temperatun. The
correlation to mo
