1

Voltammetric Determination of the Iodide/Iodine

Formal Potential and Triiodide Stability Constant

in Conventional and Ionic Liquid Media

Cameron L. Bentley,†,‡ Alan M. Bond,† Anthony F. Hollenkamp,‡,* Peter J. Mahon§,* and Jie

Zhang†

†School of Chemistry, Monash University, Clayton, Vic 3800, Australia

‡CSIRO Energy, Box 312, Clayton South, Vic 3169, Australia

§Faculty of Science, Engineering and Technology, Swinburne University of Technology,

Hawthorn, Vic 3122, Australia

2

Abstract. The iodide/triiodide/iodine (I‒/I3‒/I2)

redox system has been the subject of

electrochemical investigations for well over half a century and remains a contemporary

research interest due to the integral role of the I‒/I3‒ couple in dye-sensitized solar cell (DSSC)

technology. In this study, we have calculated the formal potential (E0’) of the I‒/I2 process and

the stability constant (Kstab) of I3‒ in two protic solvents (water and ethanol), two aprotic

solvents (acetonitrile and propylene carbonate), eight aprotic ionic liquids (AILs) and one

protic ionic liquid (PIL) using the voltammetric methodology developed herein. Furthermore,

using 1-ethyl-3-methylimidazlium bis(trifluoromethanesulfonyl)imide (abbr. [C2mim][NTf2])

as a ‘model’ ionic liquid-based DSSC electrolyte system, we have also investigated the

influence of three common additives/impurities in DSSCs (i.e., tert-butylpyridine, Li+ and

water) on the parameters E0’(I‒/I2) and Kstab and characterized two analogous redox systems,

Br‒/Br3‒/Br2 and SeCN‒/(SeCN)3

‒/(SeCN)2. E0’(I‒/I2) and Kstab(I3

‒) increase in the order ethanol

≈ acetonitrile < propylene carbonate < AILs < PIL < water; and water < ethanol ≈ PIL <

acetonitrile ≈ AILs < propylene carbonate, respectively. In the presence of the

additives/impurities (see above), E0’(I‒/I2) and Kstab increase in the order

0.5 M tert-butylpyridine < neat [C2mim][NTf2] ≈ 0.3 M Li+ < 2 wt% water and

0.5 M tert-butylpyridine << 2 wt% water < 0.3 M Li+ ≈ neat [C2mim][NTf2], respectively.

Finally, E0’(X‒/X2) and Kstab(X3‒) increase in the order SeCN‒/(SeCN)2 ≈ I‒/I2 < Br‒/Br2 and

(SeCN)3‒ << Br3

‒ < I3‒, respectively in [C2mim][NTf2]. The trends in the

(pseudo)halide/(pseudo)halogen formal potentials and tri(pseudo)halide stability constants

have been rationalized in terms of the physicochemical parameters (i.e., polarity, Gutmann

donor/acceptor numbers, ionic strength etc.) of the respective solvent/ionic liquid media.

3

Introduction

The electrochemical behavior of the iodide/triiodide/iodine (I‒/I2/I3‒) redox system has

been intensively studied for well over half a century1 and remains a contemporary research

interest due to the integral role of the I‒/I3‒ couple in dye-sensitized solar cell (DSSC)

technology. The DSSC, first reported by O’Regan and Grätzel in 19912, has been proposed as

a viable alternative for traditional (p-/n-) silicon photovoltaics in a range of applications due to

low manufacturing costs and design versatility (i.e., size, shape and flexibility).

In essence, a prototypical (n-type) DSSC has three fundamental components: (1) a TiO2

semiconductor photoanode with an adsorbed photoactive dye; (2) a platinized counter electrode

(cathode) and; (3) an electrolyte solution containing the oxidized and reduced forms of a

suitable redox couple so as to establish a redox shuttle system that regenerates the reduced form

of the dye.3, 4 As previously alluded to, the I‒/I3‒ couple was the redox mediator or ‘shuttle’

system employed in the original work.2 Since then, the corrosive and photochemical properties

of iodine have been a strong driving force for the development of alternative redox shuttle

systems, and although some promising candidates have been reported (e.g., Fc/Fc+,

SeCN/(SeCN)3‒ and Br‒/Br3

‒), the I‒/I3‒ couple still yields the most stable and efficient DSSCs.

A ‘typical’ DSSC electrolyte consists of the redox shuttle system (i.e., an iodide salt plus

iodine) plus a number of additives dissolved in pure or mixed molecular solvents (e.g.,

acetonitrile, propylene carbonate, ethanol and water). Unfortunately, the use of molecular

solvent/electrolyte media often imposes restrictions on device performance due to the poor long

term stability (high solvent volatility) and safety (solvent flammability) under light soaking

conditions.3, 4 These issues can be mitigated by employing an electrolyte solution that is based

on an appropriate ionic liquid (IL), which are typically non-volatile and display high chemical,

electrochemical and thermal stability.4-7

4

The I‒/I2 redox process has been characterized extensively in a range of conventional

solvents, most notably water8-10 and acetonitrile11-16. On inert electrode materials such as

platinum or glassy carbon, iodide is oxidized to molecular iodine in an overall one-electron per

iodide ion process:

2I− ⇌ I2 + 2e− (1)

Iodide is a Lewis base (nucleophile) and iodine is a Lewis acid (electrophile), which means

these species can combine homogeneously to form the polyhalogen complex anion, triiodide:

I− + I2 ⇌ I3− (2)

The driving force for the formation of triiodide is sensitive to donor-acceptor interactions with

the solvent17 and the equilibrium (stability) constant (i.e., Kstab) of the reaction given in Eq. 2

is therefore highly solvent dependent, ranging from ca. 103 in water to ca. 107 in acetonitrile.1

The extent to which the homogeneous process given in Eq. 2 influences the electron transfer

process given in Eq. 1 depends upon the magnitude of Kstab and the bulk concentration of I‒/I2,

and under conditions where the formation of I3‒ is favoured (i.e., large Kstab and/or high

concentrations), iodide oxidation/iodine reduction occurs in two resolved steps under

voltammetric conditions:

3I− ⇌ I3− + 2e− (3)

I3− ⇌

3

2I2 + e− (4)

The difference in the formal potentials (E0’) of the I‒/I3‒ and I3

‒/I2 processes (Eqs. 3 and 4,

respectively) is proportional to Kstab, which is a fundamentally important parameter in DSSC

applications, as it governs the energetics (thermodynamics) of the I‒/I3‒ process and influences

the amount of corrosive iodine present in the electrolyte.3

5

In previous studies18, 19, we demonstrated that the I‒/I2 redox process in the IL 1-ethyl-

3-methylimidazolium bis(trifluoromethanesulfonyl)imide is analogous to that in acetonitrile,

occurring via a triiodide intermediate. We also developed a model20 to simulate the iodide

oxidation process and estimated the stability constant of triiodide to be 106.4 in this media. We

build upon those studies here by characterizing the I‒/I2 redox process in two protic solvents

(water and ethanol), two aprotic solvents (acetonitrile and propylene carbonate), eight aprotic

ionic liquids and one protic ionic liquid (structures shown in Scheme 1). In essence, we have

developed and applied voltammetric methodology in this study to calculate the diffusion

coefficients of I‒, I3‒ and I2; the formal potentials of the I‒/I3

‒ and I3‒/I2 processes; and the

stability constant of triiodide, and then related the trends in the data to the donor/acceptor

properties of the respective solvents/ILs. We also have modelled the voltammetry of the iodide

oxidation process in each of the ILs using the experimentally determined parameters.

6

Scheme 1. Names, abbreviations and structures of the constituent cation/anions of the ILs used

in this study.

1-alkyl-3-methylimidazolium, [Cxmim]+

1-butyl-1-methylpyrrolidinium, [C4mpyr]+

N,N-diethyl-N-methyl-N-(2-

methoxyethyl)ammonium, [DEME]+

triethylammonium, [NH,2,2,2]+

bis(trifluoromethanesulfonyl)imide, [NTf2]‒

trifluoromethanesulfonate, [TfO]‒

tetrafluoroborate, [BF4]‒

hexafluorophosphate, [PF6]‒

N+

N

CH3

R

N+

CH3

CH3

CH3

N+

CH3

CH3

OCH3

N+

CH3

CH3

CH3

H

SN

- SCF3

O

O

F3C

O

O

S

O

O O-

CF3

B-

F

F

F

F

P- F

F

FF

F

F

7

Experimental Section

Reagents. 1-ethyl-3-methylimidazolium bis(trifluoromethanesulfonyl)imide

([C2mim][NTf2], Io-li-tec), 1-butyl-3-methylimidazolium bis(trifluoromethanesulfonyl)imide

([C4mim][NTf2], Solvent Innovation), 1-butyl-1-methylpyrrolidinium

bis(trifluoromethanesulfonyl)imide ([C4mpyr][NTf2], Merck), N,N-diethyl-N-methyl-N-(2-

methoxyethyl)ammonium bis(trifluoromethanesulfonyl)imide ([DEME][NTf2], Kanto

Chemical Company), 1-ethyl-3-methylimidazolium trifluoromethanesulfonate

([C2mim][OTf], Merck) and 1-butyl-3-methylimidazolium tetrafluoroborate ([C4mim][BF4],

Solvent Innovation) were commercial samples. 1-octyl-3-methylimidazolium

bis(trifluoromethanesulfonyl)imide ([C8mim][NTf2]) was prepared by a metathesis reaction

between lithium bis(trifluoromethanesulfonyl)imide (Li[NTf2], 3M Fluorad) and 1-octyl-3-

methylimidazolium chloride ([C8mim]Cl, Merck) in de-ionized water (Millipore Milli-Q Plus

185). Following preparation, [C8mim][NTf2] was taken up in dichloromethane (Merck,

EMSURE) and rinsed repeatedly with water to extract residual LiCl until the aqueous phase

passed the AgNO3 test. Triethylammonium bis(trifluoromethanesulfonyl)imide

([NH,2,2,2][NTf2]) was prepared by a metathesis reaction between Li[NTf2] and

triethylammonium chloride ([NH,2,2,2]Cl, Sigma-Aldrich, recrystallized from ethanol). 1-butyl-

3-methylimidazolium hexafluorophosphate ([C4mim][PF6]) was prepared by a metathesis

reaction between potassium hexafluorophosphate (K[PF6], Sigma-Aldrich) and 1-butyl-3-

methylimidazolium chloride ([C4mim]Cl, Solvent Innovation). Before use, each of the ILs was

dried under high vacuum (≤ 10−2 mbar) at 45°C for at least 48 hours.

1-ethyl-3-methylimidazolium iodide ([C2mim]I, Io-li-tec) was recrystallized twice

from a 2:1 mixture of ethyl acetate (Merck, EMSURE) and isopropanol (Merck, EMSURE)

and then dried under high vacuum prior to use. Tetrabutylammonium triiodide ([NBu4][I3])

was prepared by adding iodine (I2, Sigma-Aldrich) to tetrabutylammonium iodide ([NBu4]I,

8

Sigma-Aldrich) in methanol, which resulted in the precipitation of dark crystals upon mixing.

The crude product was separated, rinsed twice with methanol and then recrystallized with

ethanol, before drying at 50°C under nitrogen to a constant weight (Tm = 69 – 70°C). Care was

taken during handling and storage of [C2mim]I and [NBu4][I3] to avoid exposure to light.

Lithium nitrate (Li[NO3], Sigma-Aldrich), lithium iodide (LiI, Sigma-Aldrich),

tetrabutylammonium bromide ([NBu4]Br, Sigma-Aldrich, ≥99%), potassium selenocyanate

(K[SeCN], Sigma-Aldrich, ≥99%), ethanol (Merck, 0.01% max water), acetonitrile (Sigma-

Aldrich, anhydrous), propylene carbonate (Sigma-Aldrich, anhydrous), tert-butylpyridine

(Sigma-Aldrich, 99%), sulfuric acid (Univar), ferrocene (Fc, Sigma-Aldrich),

ferrocenemethanol (FcMeOH, Sigma-Aldrich) and silver nitrate (AgNO3, BDH, 99.9%) were

used as supplied by the manufacturer. All water/oxygen sensitive reagents were stored and

handled under a dry argon atmosphere in a glovebox.

Electrochemical systems and procedures. All voltammetric experiments were carried

out under benchtop conditions at ambient temperature (24 ± 1°C) with a Gamry Reference 600

Potentiostat/Galvanostat/ZRA (Gamry Instruments, USA). All solvents were degassed with N2

prior to experimentation and a blanket of N2 was maintained during the course of the

voltammetric experiments. A faraday cage was employed to minimize noise in all

microelectrode experiments. Positive feedback iRu compensation (Ru = uncompensated

resistance) was employed in macroelectrode experiments (Ru was estimated by electrochemical

impedance spectroscopy). All voltammetric experiments were carried out using a standard 3-

electrode arrangement with a working and reference electrode as described below and a Pt wire

auxiliary electrode. An Ag wire which had been immersed in the solution under investigation

(i.e., neat ionic liquid or solvent + supporting electrolyte) and sealed in a fritted (Vycor glass)

glass tube served as the pseudo reference electrode. In all non-aqueous electrolyte media, the

pseudo reference electrode potential was calibrated against the formal potential of the IUPAC

9

recommended Fc/Fc+ process21 in the electrolyte of interest, taking into careful consideration

the difference in the diffusion coefficients of Fc and Fc+.22, 23 In aqueous media, the pseudo

reference electrode potential was initially calibrated against the formal potential of the

FcMeOH/[FcMeOH]+ process (E0’ = 0.185 V vs. SCE)24 and later corrected to the Fc/Fc+ scale

(E0’ = 0.159 V vs. SCE)25.

The Pt macrodisk with a nominal diameter of 1.6 mm was purchased from BASi

(Bioanalytical Systems, USA) and the Pt microdisk with a nominal diameter of 20 µm was

purchased from Metrohm (Switzerland). The Pt macrodisk electrode was activated by polishing

with successively smaller (1 and 0.3 µm) aqueous alumina slurries (Kemet, UK) on a clean

polishing cloth (Buehler, USA). Adherent alumina was removed by sonication in de-ionized

water. The Pt microdisk electrode was activated by polishing with an aqueous slurry of 0.3 µm

alumina and rinsed thoroughly with de-ionized water. Prior to experimentation, the relevant

electrodes were preconditioned in 0.1 M sulphuric acid by scanning between the oxygen and

hydrogen evolution reactions26 with subsequent rinsing in de-ionized water and acetone. The

active electrode area (A) of each of the electrodes was calibrated with convolution

voltammetry20, 27, 28, using the oxidation of a Fc solution of known concentration (2.0 mM in

acetonitrile containing 0.10 M [NBu4][PF6]) and adopting a diffusion coefficient of 2.4 × 10−5

cm2 s−1, as published under these conditions.29

Data treatment, processing and simulation. The algorithm used to calculate the

convolved currents has been reported previously.28 Diffusion coefficients (D) and bulk

concentrations (Cb) were calculated simultaneously using chronoamperometry, as reported by

Compton and co-workers22, 30. Using this procedure, DI‒, DI3‒ and DI2 were calculated in each

solvent using solutions of [C2mim]I (LiI in aqueous media), [NBu4][I3] (LiI + I2 in aqueous

media) and I2, respectively.

10

Cyclic voltammetric simulations were carried out using the commercially available

DigiElch software package (v. 7F, Elchsoft, Germany) using the following mechanism, as

proposed in a previous publication20:

I2 + 2e− ⇌ I− + I− ; 𝐸0′, 𝛼, 𝑘s (5)

I2 + I− ⇌ I3− ; 𝐾stab, 𝑘f (6)

where E0, ks, α, Kstab and kf are the formal potential, standard heterogeneous electron-transfer

rate constant, transfer coefficient, triiodide stability constant and bimolecular (forward) rate

constant respectively. In all simulations, kf was arbitrarily set to 1016 M‒1 s‒1 to ensure Eq. 6 is

not limiting on the voltammetric timescale and the parameters, uncompensated resistance and

double layer capacitance were assumed to be negligible. In addition, inlaid disk electrode

geometry (r0 = 0.082 cm) was assumed in all simulations and two-dimensional (radial)

diffusion was considered. Experimentally derived D (±10%), Cb, E0’ (±10 mV) and Kstab values

were employed in the simulations, while the ks and α were systematically varied to heuristically

achieve the best fit with the experimental data over a wide range of voltammetric scan rates.

11

Results and Discussion

Electro-Oxidation of I‒ in Water, Ethanol, Acetonitrile and Propylene Carbonate.

Iodide electro-oxidation on a platinum macrodisk electrode was initially investigated in

aqueous media, where it is well-known to occur in a single step17, 31, producing molecular

iodine as per Eq. 1; representative cyclic voltammograms are shown in Figure 1a. Evidently, a

single, chemically reversible, one-electron per iodide-ion oxidation process is observed in the

potential region approximately 0.25 V positive of the Fc/Fc+ process.

Despite the fact that only a single process is observed voltammetrically (see Figure 1a),

the iodide oxidation process is complicated by a homogeneous chemical process (i.e., triiodide

formation, see Eq. 2).31 This can be illustrated by modelling the iodide oxidation process

(details are included in the Experimental Section), as is shown in Figure 1b. The dashed line in

the figure corresponds to an uncomplicated, chemically reversible oxidation process with 2:1

reactant to product (i.e., 2I‒:I2) stoichiometry, as is shown in Eq. 5. The dotted line corresponds

to the reaction shown in Eq. 5, coupled to a bimolecular homogeneous process between the

reactant and product, as is shown in Eq. 6. The diffusion coefficients were determined

experimentally (see Experimental Section), Kstab was taken to be 102.9 as reported elsewhere1

and E0’, α and ks were adjusted to achieve the best fit with the experimental data (solid line in

Figure 1b). Clearly, the dashed line (Eq. 5) is not in agreement with the experimental data,

featuring much sharper oxidation and reduction peaks. The dotted line (Eqs. 5 and 6) on the

other hand, is in good agreement with the experimental data over scan rates ranging from at

least 10 mV s‒1 to 100 mV s‒1 (see Figure S1), reinforcing the fact that triiodide formation

influences the iodide oxidation process in aqueous media.

Further investigations on the iodide oxidation process were carried out in acetonitrile,

propylene carbonate and ethanol; representative normalized cyclic voltammograms are shown

in Figure 2. In acetonitrile, two chemically reversible processes separated by approximately

12

0.4 V are observed, corresponding to the I‒/I3‒ (see Eq. 3) and I3

‒/I2 (see Eq. 4) processes at

lower and higher potentials respectively.14, 31 As previously discussed, the potential gap

separating the two processes is proportional to Kstab, which is consistent with the fact that the

stability constant of triiodide is approximately four orders-of-magnitude higher in acetonitrile

(Kstab ≈ 107) compared to water.1 Similar iodide oxidation behavior also is observed in

propylene carbonate, where again, two chemically reversible processes separated by

approximately 0.5 V are evident, consistent with the relatively high triiodide stability constant

in this solvent (Kstab ≈ 108).32 Finally, the cyclic voltammetric response of iodide in ethanol is

intermediate between that observed in water and acetonitrile: two overlapping, chemically

reversible processes separated by approximately 0.2 V are observed, which is consistent with

the reported stability constant (Kstab ≈ 105) of triiodide in this media.33

13

Figure 1. (a) Cyclic voltammograms obtained from 2.5 mM LiI in aqueous media (0.2 M

Li[NO3]) at a 1.6 mm dia. Pt macrodisk electrode with scan rates (from top to bottom) of 100,

50, 25 and 10 mV s‒1. (b) Comparison of experimental (50 mV s‒1, solid line) and simulated

data obtained using the mechanisms described by Eq. 5 (dashed line) or Eqs. 5 and 6 (dotted

line). The following parameters were used in the simulations: E0’ = 0.223 V, α = 0.5, ks = 1

cm/s, Keq = 102.9, DI‒ = 1.6 × 10‒5 cm2 s‒1, DI2 = 1.3 × 10‒5 cm2 s‒1, DI3‒ = 1.1 × 10‒5 cm2 s‒1.

-16

-12

-8

-4

0

4

8

12

16

20

0 0.1 0.2 0.3 0.4 0.5

I(µ

A)

E (V) vs. Fc/Fc+

(a)

-10

-5

0

5

10

0 0.1 0.2 0.3 0.4 0.5

I(µ

A)

E (V) vs. Fc/Fc+

(b)

14

Figure 2. Cyclic voltammograms (normalized to the I‒/I3‒ oxidation peak current) obtained

from the electro-oxidation of (from top to bottom) 2.5 mM LiI in water (+0.2 M Li[NO3]), 2.7

mM [C2mim]I in acetonitrile (+0.2 M [C2mim][NTf2]), 4.5 mM [C2mim]I in propylene

carbonate (+0.2 M [C2mim][NTf2]) and 2.6 mM [C2mim]I in ethanol (+0.2 M [C2mim][NTf2])

at a 1.6 mm dia. Pt macrodisk electrode with a scan rate of 100 mV s‒1. The arrows indicate

zero current for each of the cyclic voltammograms.

-0.8 -0.6 -0.4 -0.2 0 0.2 0.4 0.6I

/ I p

E (V) vs. Fc/Fc+

Ethanol, Kstab = 104.8

Propylene Carbonate, Kstab = 107.8

Acetonitrile, Kstab = 107.1

Water, Kstab = 102.9

15

Calculation of E0’(I‒/I2) and Kstab. The Nernst expression29 for the overall I‒/I2 process

(see Eq. 1) is:

𝐸 = 𝐸0′(I−/I2) +

𝑅𝑇

2𝐹ln (

[I2]

[I−]2) (7)

where R is the universal gas constant, T is temperature and F is Faraday’s constant. Due to the

absence of information relating to the activity of solutes, particularly in ionic liquid media34,

concentrations have been used in place of activities in Eq. 7 and all subsequent equations.

Under conditions where mass transport is governed solely by semi-infinite planar diffusion

(i.e., at a macrodisk electrode) and I‒ is the only species initially present in solution, the

following relationship between E0’(I‒/I2) and the reversible half-wave potential, E1/2, can be

derived using the diffusion layer method29, 35-37:

𝐸0′(I−/I2) = 𝐸1/2(I−/I2) −

𝑅𝑇

2𝐹[ln (

√𝐷I−

√𝐷I2

) − ln[I−]b] (8)

where the subscript ‘b’ signifies ‘bulk concentration’. E1/2 is the potential value corresponding

to half the limiting current (I1/2) obtained from a steady-state voltammogram for a reversible

process, and can be readily estimated from a transient cyclic voltammogram as follows:

𝐸1/2 ≅𝐸p,ox + 𝐸p,red

2= 𝐸mid (9)

Where Ep,ox and Ep,red are the oxidation and reduction peak potentials, respectively. When a

‘split’ iodide oxidation response is observed (i.e., in all of the solvents investigated above

except water), E1/2(I‒/I2) cannot be directly estimated from the cyclic voltammogram, however

it can be easily calculated from the E1/2(I‒/I3

‒) and E1/2(I3‒/I2) values, since:

𝐸1/2(I−/I2) =2𝐸1/2(I−/I3

−) + 𝐸1/2(I3−/I2)

3 (10)

16

Using this method with DI‒ and DI2 being determined by chronoamperometry, as outlined in

the Experimental Section, E0’(I‒/I2) was calculated to be ‒0.11 V vs. Fc/Fc+ in acetonitrile (see

Figure 2), which is in excellent agreement with the value reported in the literature.1

The Nernst expression29 for the I‒/I3‒ process (see Eq. 3) is:

𝐸 = 𝐸0′(I−/I3

−) +𝑅𝑇

2𝐹ln (

[I3−]

[I−]3) (11)

Again, from the diffusion layer method29, 35-37, it follows that:

𝐸0′(I−/I3

−) = 𝐸1/2(I−/I3−) −

𝑅𝑇

2𝐹[ln (

√𝐷I−

√𝐷I3−

) − 2ln[I−]b + ln (4

3)] (12)

From this equation (DI‒ and DI3‒ were again calculated using chronoamperometry, as outlined

in the Experimental Section), E0’(I‒/I3‒) was calculated to be ‒0.33 V vs. Fc/Fc+ in acetonitrile,

which is also in excellent agreement with literature reports.1

In order to calculate Kstab, Eq. 11 can be subtracted from Eq. 7 to give:

𝐸0′(I−/I2) − 𝐸0′

(I−/I3−) =

𝑅𝑇

2𝐹ln(𝐾stab) (13)

Substituting the appropriate values into Eq. 13, Kstab is calculated to be 107.1 in acetonitrile,

which once again is in excellent agreement with what has been reported in the literature.1

Finally, if Eq. 13 is combined with Eqs. 8, 10 and 12, Eq. 14 can be derived, assuming mass

transport is governed solely by semi-infinite planar diffusion and I‒ is the only species initially

present in solution:

ln(𝐾𝑠𝑡𝑎𝑏) =2𝐹

3𝑅𝑇[𝐸1/2(I3

−/I2) − 𝐸1/2(I−/I3−)] + ln (

4√𝐷I2

3√𝐷I3−[I−]b

) (14)

Thus, using Eq. 14, Kstab can be readily calculated from the cyclic voltammetric response of

iodide at a macrodisk electrode, which gives E1/2(I3‒/I2) and E1/2(I

‒/I3‒), provided DI2, DI3

‒ and

17

[I‒]b are known. The E0’(I‒/I2), E0’(I‒/I3

‒) and Kstab values calculated using Eqs. 8, 12 and 14,

respectively, in the molecular solvents of interest in this study are summarized in Table 1.

Table 1. DI‒, DI3‒, DI2

, E0’(I‒/I2), E0’(I‒/I3

‒) and Kstab values calculated using voltammetry in

water, acetonitrile, propylene carbonate and ethanol. Literature values available are provided

in parenthesis.

Solvent/Supporting

Electrolyte

DI‒ / 10‒5

cm2 s‒1

DI3‒ / 10‒5

cm2 s‒1

DI2 / 10‒5

cm2 s‒1

E0’(I‒/I2) / V

vs. Fc/Fc+ E0’(I‒/I3

‒) / V

vs. Fc/Fc+ log10(Kstab)

Water / 0.2 M LiNO3 1.6 1.1 1.3 0.22a

(0.221)1

0.14b

(0.136)1

2.9b (2.9)1

Acetonitrile / 0.2 M

[C2mim][NTf2] 2.1 2.1 2.4

‒0.11

(‒0.12)1

‒0.32

(‒0.33)1 7.1 (7.4, 6.6)1

Propylene Carbonate /

0.2 M [C2mim][NTf2] 0.35 0.41 0.50 ‒0.081 ‒0.31 7.8 (7.8)38

Ethanol / 0.2 M

[C2mim][NTf2] 0.60 0.73 0.81 ‒0.12 ‒0.26 4.8 (4.7)33

aInitially calculated vs. FcMeOH/[FcMeOH]+ and converted to the Fc/Fc+ scale as described in the Experimental Section. bEstimated using a numerical simulation

As shown in Table 1, the values of parameters calculated in this work are in excellent

agreement with those reported in the literature (where available). The diffusion coefficients for

I‒, I3‒ and I2 increase in the order propylene carbonate < ethanol < water < acetonitrile, while

viscosities increase in the inverse order39, 40, in accordance with predictions based on the

Stokes-Einstein Relation.41 Interestingly, the ratio, DI‒/DI2 varies markedly, ranging from 0.7

in propylene carbonate to 1.2 in water, which is probably attributable to relative differences in

the solvation of I‒ and I2 in each solvent. As discussed below, this effect is much more

pronounced in IL media.

Assuming that the Fc/Fc+ formal potential is solvent independent21, E0’(I‒/I2) is

significantly more positive in water compared to any of the other molecular solvents

investigated in this work. In other words, the oxidation of iodide is most difficult in aqueous

media, which is not surprising given that the high polarity and the large dielectric constant (εr

18

= 78.4)42 of water favors the formation of ions (i.e., I‒) over neutral molecules (i.e., I2). In

addition, water is the strongest Lewis acid (electrophile) of any of the solvents12, 17, evidenced

by its large Gutmann acceptor number (AN = 54.8)42, meaning it can interact with and stabilize

iodide, a Lewis base (nucleophile), to the greatest extent (i.e., via strong donor/acceptor

interactions). Although water is also a good Lewis base (Gutmann donor number, DN = 18.0

kcal mol‒1)42, it is not expected to interact with non-polar iodine (a Lewis acid) to any

significant extent, which is consistent with the fact that iodine is sparingly soluble in aqueous

media ([I2]SAT = 1.2 mM).9 Ethanol, acetonitrile and propylene carbonate are less polar and

possess lower dielectric constants than water (εr = 24.3, 36.0 and 64.4, respectively)39, 42 and

would therefore be expected to interact with charged iodide and uncharged iodine to lesser and

greater extents, respectively, explaining the relatively more negative E0’(I‒/I2) values in these

solvents. This is consistent with the increased solubility of iodine in these solvents, for

example, at 25°C iodine is soluble up to 21.4 wt% (4.7 mol%) in ethanol.43

The driving force for the formation of triiodide (i.e., Kstab) is governed by how strongly

the solvent interacts with or ‘solvates’ (i.e., by donor-acceptor type interactions) the three

species shown in Eq. 2, iodide, iodine and/or triiodide. Strong stabilizing interactions, such as

that between water and charged species (i.e., I‒, discussed above) are expected to decrease the

driving force for triiodide formation, which explains why Kstab is lowest (Kstab = 102.9) in

aqueous media. Relative to water, the aprotic solvents, acetonitrile and propylene carbonate,

are weakly solvating, with comparably low Gutmann donor (DN = 14.1 and 15.1 kcal mol‒1,

respectively)39, 42 and acceptor (AN = 18.9 and 18.3, respectively)39, 42 numbers, explaining

why the driving force for triiodide formation is much higher (Kstab = 107.1 and 107.8,

respectively) in these solvents. Finally, despite possessing the lowest dielectric constant,

ethanol possesses a Gutmann donor number greater than that of water (DN = 20.0 kcal mol‒

19

1)42 and an acceptor number which is larger than those of the aprotic solvents (AN = 37.1)42,

explaining the intermediate driving force (Kstab = 104.8) for triiodide formation in this solvent.

20

The I‒/I3‒/I2 Redox System in Ionic Liquid Media. The iodide/triiodide/iodine redox

system was characterized electrochemically in a range of ILs; normalized cyclic

voltammograms obtained from the electro-oxidation of iodide are shown in Figure 3. In all of

the ILs, iodide oxidation clearly occurs in two steps, attributable to the I‒/I3‒ and I3

‒/I2 processes

at lower and higher potentials, respectively, and in agreement with previous reports.18, 30, 44 In

the ILs which do not contain dissociable protons (termed aprotic ionic liquids, AILs), the

potential gap separating the I‒/I3‒ and I3

‒/I2 processes is comparable to that in the aprotic

solvents acetonitrile and propylene carbonate (see Figure 2), qualitatively indicating that Kstab

is comparable in these physicochemically disparate media. By contrast, the potential gap

separating the I‒/I3‒ and I3

‒/I2 processes in IL which contains a dissociable proton (termed a

protic ionic liquid, PIL), [NH,2,2,2][NTf2] (see Figure 3, top) is comparable to that in the protic

solvent ethanol (see Figure 2, bottom). E0’(I‒/I2), E0’(I‒/I3

‒) and Kstab were quantified in each of

the ILs using Eqs. 8, 12 and 14, respectively; the results are summarized in Table 2.

Table 2. DI‒, DI3‒, DI2, E

0’(I‒/I2), E0’(I‒/I3

‒) and Kstab values calculated using voltammetry in a

range of ionic liquids. Room temperature (298 K) viscosity data obtained from the literature

also is included in the table.

Ionic Liquid η (cP) DI‒ / 10‒7

cm2 s‒1

DI3‒ / 10‒7

cm2 s‒1

DI2 / 10‒7

cm2 s‒1

E0’(I‒/I2) / V

vs. Fc/Fc+ E0’(I‒/I3

‒) / V

vs. Fc/Fc+ log10(Kstab)

[NH,2,2,2][NTf2] 5445 2.0 2.3 6.5 0.10 ‒0.037 4.7

[DEME][NTf2] 6846 0.81 1.7 5.2 ‒0.034 ‒0.23 6.7

[C4mpyr][NTf2] 7647 0.84 1.7 5.5 ‒0.035 ‒0.23 6.7

[C8mim][NTf2] 9248 0.72 1.4 5.7 ‒0.063 ‒0.27 6.9

[C4mim][NTf2] 5548 1.2 2.3 7.8 ‒0.032 ‒0.23 6.6

[C2mim][NTf2] 3248 2.4 4.4 11 ‒0.015 ‒0.21 6.5

[C2mim][OTf] 5149 2.1 2.6 6.7 ‒0.010 ‒0.19 6.1

[C4mim][BF4] 10050 0.93 1.6 4.6 0.017 ‒0.19 6.9

[C4mim][PF6] 26150 0.53 0.66 1.8 0.016 ‒0.19 7.1

21

Figure 3. Cyclic voltammograms (normalized to the I‒/I3‒ oxidation peak current) obtained

from the electro-oxidation of I‒ in (from top to bottom) [NH,2,2,2][NTf2], [DEME][NTf2],

[C4mpyr][NTf2], [C2mim][NTf2], [C4mim][NTf2], [C8mim][NTf2], [C2mim][OTf],

[C4mim][BF4] and [C4mim][PF6] at a 1.6 mm dia. Pt macrodisk electrode with a scan rate of

50 mV s‒1. The concentration of [C2mim]I in each of the ILs was (from top to bottom) 6.2 mM,

13.9 mM, 12.2 mM, 8.5 mM, 6.2 mM, 12.4 mM, 9.9 mM, 13.4 mM and 14.9 mM. The arrows

indicate zero current for each of the cyclic voltammograms.

-0.8 -0.6 -0.4 -0.2 0 0.2 0.4 0.6 0.8I

/ I P

E (V) vs. Fc/Fc+

[C4mim][PF6], Kstab = 107.1

[C4mim][BF4], Kstab = 107.0

[C2mim][OTf], Kstab = 106.1

[C8mim][NTf2], Kstab = 106.9

[C4mim][NTf2], Kstab = 106.6

[C2mim][NTf2], Kstab = 106.5

[C4mpyr][NTf2], Kstab = 106.7

[DEME][NTf2], Kstab = 106.7

[NH,2,2,2][NTf2], Kstab = 104.7

22

The diffusivities of iodide, triiodide and iodine all generally adhere to the Stokes-

Einstein relation41, increasing in the order [C4mim][PF6] < [C4mim][BF4] ≈ [C8mim][NTf2] ≈

[C4mpyr][NTf2] ≈ [DEME][NTf2] < [C4mim][NTf2] ≈ [NH,2,2,2][NTf2] ≈ [C2mim][OTf] <

[C2mim][NTf2], in accordance with the relative fluidities (i.e., 1/η) of the ILs. In all cases,

diffusivities increase in the order I‒ < I3‒ < I2, which is likely attributable to the relative charge

density of these species. In other words, electrostatic interactions with the ionic medium

hinders the mass-transport of I‒ to a greater extent than I3‒ (diffuse negative charge) or I2

(uncharged). The ratio, DI‒/DI2, also varies markedly in the investigated ILs, ranging from 0.13

in [C8mim][NTf2] to 0.31 in [C2mim][OTf], which is understandable, given that the strength

of the interaction between iodide/iodine and the IL will undoubtedly be dependent on the

constituent cations and anions. Similar observations have been previously been reported for

the ferrocene/ferrocenium and cobaltocene/cobaltocenium redox couples in IL media.22

E0’(I‒/I2) is more positive in IL media compared to the nonaqueous solvents (i.e.,

acetonitrile, ethanol and propylene carbonate, see Table 1). This indicates that relative to the

nonaqueous solvents, the ILs interact with (and stabilize by donor/acceptor type interactions)

I‒ and I2 to greater and lesser extents, respectively, which is not surprising, given the high ionic

strength of this class of medium. Indeed, I2 is sparingly soluble in IL media51 and weakly

bound, as it can be easily removed (volatilized) by standard degassing protocols (i.e., vacuum

or bubbling an inert gas such as nitrogen through the media). In all of the AILs, E0’(I‒/I2) lies

in the narrow potential range ‒0.06 to 0.02 V vs. Fc/Fc+. Comparing the [NTf2]‒ containing ILs

with aprotic cations (i.e., [Cxmim]+, [C4mpyr]+ and [DEME]+, see Scheme 1), it is clear that

the structure of the cation has a very minor influence on E0’(I‒/I2). The only notable trend is

that increasing the size of the alkyl chain substituent on the imidazolium cation results in a

minor shift in E0’(I‒/I2) towards more negative potentials. Comparing the [Cxmim]+ containing

ILs (i.e., [NTf2]‒, [OTf]‒, [BF4]

‒ and [PF6]‒, see Scheme 1), it is also evident that the structure

23

of the anion also has a very minor influence on E0’(I‒/I2). Overall, the largest shift in E0’(I‒/I2)

was induced by changing from an aprotic cation to a protic cation (i.e., compare [Cxmim][NTf2]

to [NH,2,2,2][NTf2], see Scheme 1). This significant shift in E0’(I‒/I2) towards more positive

potentials is likely attributable to a (relatively) strong, stabilizing interaction between [NH,2,2,2]+

and I‒ (see below).

As discussed above, the driving force for the formation of triiodide (i.e., Kstab) is

governed by how strongly the solvent interacts with or ‘solvates’ iodide and/or iodine (see Eq.

2). In the AILs, Kstab is comparable to that in acetonitrile (see Table 1), ranging from 106.1 in

[C2mim][OTf] to 107.1 in [C4mim][PF6]. The cation structure has minimal influence on Kstab in

the AILs containing the [NTf2]‒ anion, with values of 106.6, 106.7 and 106.7 for the [C4mim]+,

[C4mpyr]+ and [DEME]+ cations, respectively. Increasing the length of the alkyl chain

substituent on the imidazolium cation increases Kstab slightly, with values of 106.5, 106.6 and

106.9 in [C2mim][NTf2], [C4mim][NTf2] and [C8mim][NTf2], respectively. This trend is

consistent with the fact that the donor and acceptor numbers (and hence the coordinating

ability) of ILs containing the [Cxmim]+ cation decrease with increasing alkyl chain length.52

Changing the structure of the anion has a more significant influence on Kstab, with values of

106.1, 106.5, 106.9 and 107.1 in [C2mim][OTf], [C2mim][NTf2], [C4mim][BF4] and [C4mim][PF6],

respectively. The increase in Kstab coincides perfectly with decreasing coordinating ability of

the ILs, with the donor numbers increasing in the order [PF6]‒ < [BF4]

‒ < [NTf2]‒ < [OTf]‒ < I‒

when paired with [Cxmim]+ cations. In addition, acceptor numbers also increase in the order

[NTf2]‒ < [BF4]

‒ < [OTf]‒ when paired with [C2mim]+.52 Finally, the PIL [NH,2,2,2][NTf2] is

again the outlier, with a Kstab value of 104.7, resembling ethanol rather than acetonitrile (see

Table 1). Although there are no donor/acceptor number data available for [NH,2,2,2][NTf2], we

attribute the two order-of-magnitude decrease in Kstab to a (relatively) strong, stabilizing

interaction between [NH,2,2,2]+ and I‒.

24

From the DSSC technology standpoint, a high triiodide stability constant is preferable,

as it minimizes the amount of free (corrosive) iodine present in the electrolyte. Even taking the

lowest Kstab value from Table 2 (Kstab = 104.7), in a typical DSSC electrolyte mixture containing

at least a 10-fold excess of iodide over iodine, the concentration of free I2 is almost negligible

(e.g., [I2]free ≈ 1 × 10‒7 M when [I2] = 0.05 M and [I‒] = 0.5 M). The operating voltage generated

by a DSSC under illumination corresponds to the difference in the Fermi potential of the

semiconductor anode (TiO2) and the redox potential of the active shuttle (i.e., I‒/I3‒).3, 4 As

shown in Eq. 13, maximizing E0’(I‒/I2) and minimizing Kstab will maximize E0’(I‒/I3‒).

Therefore, from an energetics standpoint, [NH,2,2,2][NTf2] possesses the most favourable

properties, as E0’(I‒/I3‒) is approximately 0.2 V more positive than any of its aprotic

counterparts. This comparison should be taken cum grano salis, however, as it naively assumes

the Fermi level of the semiconductor is solvent (electrolyte) independent. In reality, the

electrolyte solvent must be carefully optimized to ensure it is; (i) compatible with the redox

shuttle, dye, semi-conductor and various additives (discussed below) commonly included in

DSSC electrolytes; (ii) relatively fluid to minimize mass-transport limitations and; (iii) stable

under light-soaking conditions over an extended period of time.3, 4

25

Modelling of the electro-oxidation of I‒ in ionic liquid media. In a previous study20,

we developed and applied a relatively simple termolecular electrode reaction mechanism (see

the Experimental Section) to model the electro-oxidation of iodide in [C2mim][NTf2]. Here,

we have successfully applied this mechanism to model the electro-oxidation of iodide in a

range of ILs; representative cyclic voltammograms obtained in [C2mim][NTf2] and

[NH,2,2,2][NTf2] are shown in Figure 4. In addition, experiment-simulation comparisons for all

of the ILs investigated in this work are available in the Supporting Information (Figures S2 to

S8). In general, there is excellent agreement between the experimental cyclic voltammograms

and the simulations obtained based on the parameters outlined in Table 3, adding further

confidence that the proposed model is valid in at least the phenomenological sense. Evidently,

the D, E0’ and Kstab values derived from the simulations are in excellent agreement with those

calculated experimentally (see Table 2). As addressed in our previous publication20, due to the

unrealistically high kf value used in the simulations (see Eq. 6), the ks and α values derived

using this mechanism are not likely to be unique or quantitatively meaningful and for this

reason, no attempt will be made to interpret these kinetic parameters.

Table 3. Data extracted from the comparison of experimental and simulated cyclic

voltammetric data obtained based on the mechanism described by Eqs. 5 and 6.

Ionic Liquid DI‒ / 107 cm2

s‒1

DI3‒ / 107 cm2

s‒1

DI2 / 107 cm2

s‒1 E0’(I‒/I2) / V

ks / cm

s‒1 α log10(Kstab)

[NH,2,2,2][NTf2] 2.1 2.5 6.5 0.11 10 0.5 4.7

[DEME][NTf2] 0.78 1.6 5.2 -0.027 2.3 0.38 6.7

[C4mpyr][NTf2] 0.82 1.6 5.4 ‒0.028 1.4 0.37 6.7

[C2mim][NTf2] 2.4 4.4 11 ‒0.012 4 0.4 6.5

[C4mim][NTf2] 1.1 2.3 7.8 ‒0.026 1 0.36 6.6

[C8mim][NTf2] 0.72 1.6 5.7 ‒0.059 0.9 0.32 6.9

[C2mim][OTf] 2.1 2.8 6.9 ‒0.002 1.1 0.35 6.1

[C4mim][BF4] 0.90 1.6 4.6 0.021 1.8 0.35 6.9

[C4mim][PF6] 0.55 0.70 1.8 0.020 0.7 0.34 7.1

26

Figure 4. Comparison of the simulated (○) and experimental (—) cyclic voltammograms

obtained from the electro-oxidation of I‒ in (a) [C2mim][NTf2] ([I‒] = 8.5 mM) and (b)

[NH,2,2,2][NTf2] ([I‒] = 6.2 mM) at a 1.6 mm dia. Pt macrodisk electrode with scan rates of 10,

50 and 100 mV s‒1. Simulation parameters are available in Table 3.

-6

-4

-2

0

2

4

6

-0.5 -0.25 0 0.25 0.5 0.75

I(µ

A)

E (V) vs. Fc/Fc+

(a)

-4

-3

-2

-1

0

1

2

3

4

-0.5 -0.25 0 0.25 0.5 0.75

I(µ

A)

E (V) vs. Fc/Fc+

(b)

27

Effects of Li+, t-butylpyridine and water on E0’(I‒/I2) and Kstab in Ionic Liquid

Media. The effect that two common DSSC additives4, lithium ions (Li+) and t-butylpyridine

(t-BPy), and one common impurity in ILs, water53, has on the stability constant of triiodide in

[C2mim][NTf2] was investigated using cyclic voltammetry (shown in Figure 5). Compared to

neat [C2mim][NTf2], the cyclic voltammogram obtained in the presence of t-BPy (0.5 M, see

Figure 5, top-middle) is shifted to more negative potentials and the potential gap separating the

I‒/I3‒ and I3

‒/I2 processes (see Eqs. 3 and 4) is smaller, implying that Kstab has decreased

(quantified below). Aside from the magnitude of the measured currents, the cyclic

voltammogram obtained in the presence of Li+ (0.3 M, see Figure 5, bottom-middle) is barely

distinguishable from that obtained in neat [C2mim][NTf2], implying that this additive has

minimal influence on the I‒/I3‒/I2 redox processes. The addition of water (≈2 wt%, see Figure

5, bottom) to [C2mim][NTf2] shifts the iodide oxidation wave towards more positive potentials

and decreases the potential gap separating the I‒/I3‒ and I3

‒/I2 processes marginally, as expected

based on the discussions above. From the cyclic voltammetric data shown in Figure 5, and

assuming that the ratios DI‒/DI2 and DI2/DI3

‒ are unchanged from neat [C2mim][NTf2] (see Table

2), the values E0’(I‒/I2) and Kstab in the presence of the additives were calculated using Eqs. 8

and 14, respectively. The data are summarized below in Table 4.

Table 4. DI‒, E0’(I‒/I2), E0’(I‒/I3‒) and Kstab values calculated from voltammetric data in

[C2mim][NTf2] in the presence of t-BPy, Li+ or H2O.

Ionic Liquid Additive/Impurity DI‒ / 107

cm2 s‒1

E0’(I‒/I2) / V

vs. Fc/Fc+ E0’(I‒/I3

‒) / V

vs. Fc/Fc+ log10(Kstab)

[C2mim][NTf2] ‒ 2.5 ‒0.015 ‒0.21 6.5

[C2mim][NTf2] 0.5 M t-BPy 2.5 ‒0.087 ‒0.22 4.6

[C2mim][NTf2] 0.3 M Li[NTf2] 1.4 ‒0.013 ‒0.20 6.3

[C2mim][NTf2] ≈2 wt% H2O 5.3 0.030 ‒0.14 5.8

28

Figure 5. Concentration-normalized cyclic voltammograms obtained from the electro-

oxidation of I‒ in [C2mim][NTf2] containing (from top to bottom) no additives (neat),

0.5 M t-BPy, 0.3 M Li[NTf2] and 2 wt% H2O at a 1.6 mm dia. Pt macrodisk electrode with a

scan rate of 50 mV s‒1. The concentration of [C2mim]I in each of the ILs was (from top to

bottom) 8.5 mM, 9.4 mM, 14.7 mM and 7.7 mM. The arrows indicate zero current for each of

the cyclic voltammograms.

-0.8 -0.6 -0.4 -0.2 0 0.2 0.4 0.6 0.8I

/ C

(µ

A M

‒1)

E (V) vs. Fc/Fc+

+2 wt% H2O, Kstab = 105.8

+0.3 M Li[NTf2], Kstab = 106.3

+0.5 M t-BPy, Kstab = 104.6

neat [C2mim][NTf2], Kstab = 106.5

400 µA M‒1

29

The presence of t-BPy in [C2mim][NTf2] shifts E0’(I‒/I2) by approximately ‒0.07 V and

decreases Kstab by approximately two orders of magnitude. Analogous to its parent compound

pyridine (DN = 33.1 kcal mol‒1)42, which is well-known to form charge-transfer complexes

with iodine54, t-BPy is a strong Lewis base and is therefore expected to interact with

electrophilic iodine to a significant extent. The stabilizing interaction between iodine and

t-BPy makes iodide oxidation thermodynamically easier (i.e., negative E0’ shift) and reduces

the driving force for the formation of triiodide (i.e., decreased Kstab). The shift in E0’(I‒/I2) is

offset by the decrease in Kstab, so E0’(I‒/I3‒) remains relatively constant. The presence of Li+ in

[C2mim][NTf2] reduces DI‒ , and marginally decreases Kstab. It is widely reported55 that

dissolving Li+ into ILs increases the medium viscosity, explaining the decrease in

DI‒ (as per the Stokes-Einstein relation41), while the slight decrease in Kstab likely arises due to

a weak stabilizing interaction between Li+ and I‒. The presence of H2O in [C2mim][NTf2]

increases DI‒, shifts E0’(I‒/I2) approximately 45 mV and decreases Kstab by almost an order of

magnitude, which is unsurprising, given that DI‒, E0’(I‒/I2) and Kstab are 1.6 ×10‒5 cm2 s‒1, 0.20

V vs. Fc/Fc+ and 102.9, respectively in aqueous media (see Table 1). As discussed above, the

shifts in E0’(I‒/I2) and Kstab are likely attributable to a (relatively) strong, stabilizing interaction

between H2O and I‒.

The Br‒/Br3‒/Br2 and SeCN‒/(SeCN)3

‒/(SeCN)2 Redox Systems in Ionic Liquid

Media. In order to emphasise the general applicability of the theory developed above, the

electro-oxidation of another halide, bromide and the pseudohalide, selenocyanate were

investigated in [C2mim][NTf2]; representative cyclic voltammograms are shown in Figure 6.

Both Br‒ and SeCN‒ have been employed as redox mediators in DSSCs4 and known to undergo

the same general electrode reaction mechanism as I‒ in IL media56, 57:

30

3X− ⇌ X3− + 2e− (15)

X3− ⇌

3

2X2 + e− (16)

where X‒, X3‒ and X2 represent a general (pseudo)halide, tri(pseudo)halide and

(pseudo)halogen, respectively. Br‒ oxidation (see Figure 6, middle) occurs in the potential

region approximately 0.4 V positive of I‒ oxidation (see Figure 6, top) and the potential gap

separating the Br‒/Br3‒ and Br3

‒/Br2 processes (≈0.3 V) is marginally smaller than that

separating the I‒/I3‒ and I3

‒/I2 processes (≈0.35 V), qualitatively indicating that the stability

constant of Br3‒ is smaller than that of I3

‒. SeCN‒ oxidation (see Figure 6, bottom) occurs in

the same potential region as I‒ oxidation (see Figure 6, top), however the SeCN‒/(SeCN)3‒ and

(SeCN)3‒/(SeCN)2 processes overlap significantly, indicating that the stability constant of

(SeCN)3‒ is much smaller than that of I3

‒ and Br3‒. E0’(X‒/X2), E0’(X‒/X3

‒) and Kstab were

quantified for each of the redox systems using Eqs. 8, 12 and 14, respectively; the results are

summarized in Table 5.

Table 5. DX‒, DX3‒, DX2

, E0’(X‒/X2), E0’(X‒/X3

‒) and Kstab values calculated using voltammetry

for the I‒/I3‒/I2, Br‒/Br3

‒/Br2 and SeCN‒/(SeCN)3‒/(SeCN)2 redox systems in [C2mim][NTf2].

Redox System DX‒ / 10‒7

cm2 s‒1

DX3‒ / 10‒7

cm2 s‒1

DX2 / 10‒7

cm2 s‒1

E0’(X‒/X2) /

V vs. Fc/Fc+ E0’(X‒/X3

‒) /

V vs. Fc/Fc+ log10(Kstab)

I‒/I3‒/I2 2.4 4.4 11 ‒0.015 ‒0.21 6.5

Br‒/Br3‒/Br2 2.0 4.5 (±0.5) 7.0 (±0.7) 0.37 0.22 5.3

SeCN‒/(SeCN)3‒/(SeCN)2 3.0 3.5 (±0.4) 4.0 (±0.4) ‒0.064 ‒0.18 3.9

31

Figure 6. Concentration-normalized cyclic voltammograms obtained from the electro-

oxidation of (from top to bottom) 8.5 mM [C2mim]I, 12.9 mM [NBu4]Br and 11.8 mM

K[SeCN] in [C2mim][NTf2] at a 1.6 mm dia. Pt macrodisk electrode with a scan rate of 50 mV

s‒1. The arrows indicate zero current for each of the cyclic voltammograms.

-0.8 -0.6 -0.4 -0.2 0 0.2 0.4 0.6 0.8 1I

/ C

(µ

A M

‒1)

E (V) vs. Fc/Fc+

I‒/I2, E0' = ‒0.015 V vs. Fc/Fc+, Kstab = 106.5

200 µA M‒1

Br‒/Br2, E0' = 0.37 V vs. Fc/Fc+, Kstab = 105.3

SeCN‒/(SeCN)2, E0' = ‒0.064 V vs. Fc/Fc+, Kstab = 103.9

32

The diffusivities of Br3‒/Br2 and (SeCN)3

‒/(SeCN)2 were estimated indirectly using

double-step chronoamperometry at a microdisk electrode (see Figures S9 to S12) in solutions

of [NBu4]Br and K[SeCN], respectively, as reported by Compton and co-workers.58 As

indicated in Table 5, there is significantly more uncertainty (± 10%) associated with the

outcome of fitting of the data in this way than if the values were determined directly (as

performed for the I‒/I3‒/I2 redox system, see the Experimental Section).

As previously alluded to, Br‒ is thermodynamically harder the oxidize than I‒, with

E0’(Br‒/Br2) lying approximately 0.39 V positive of E0’(I‒/I2). Again, because E0’(Br‒/Br2) and

E0’(I‒/I2) are solvent dependent, the difference between them also varies from solvent to

solvent, with values of 0.35 and 0.47 in acetonitrile and water, respectively.1 The stability

constant of Br‒ is approximately one order of magnitude lower than that of I3‒ in

[C2mim][NTf2]. At first glance, the order of the trihalide stability constants appears to be

incorrect, since Br‒ is a stronger Lewis base than I‒ and Br2 is a stronger Lewis acid than I2.

Indeed, Kstab(Br3‒) is marginally larger than Kstab(I3

‒) in the aprotic solvents acetonitrile,

nitromethane and acetone.17 We attribute the seemingly ‘reversed’ order of Kstab values to the

stronger donor/acceptor or electrostatic interactions that likely exist between the ionic medium

and Br‒/Br2 couple (particularly Br‒) compared to the more charge-diffuse I‒/I2 couple.

Interestingly, in aqueous media, the stability constant of trihalide species also increases in the

order Cl3‒ < Br3

‒ < I3‒, due to the fact that hydration (solvation) of the X‒/X2 species increases

in the order I‒/I2 < Br‒/Br2 < Cl‒/Cl2.

SeCN‒ is slightly easier to oxidize than I‒, with E0’(X‒/X2) being approximately 0.05 V

more negative for the former species. The stability constant of (SeCN)3‒ is almost three orders

of magnitude lower than that of I3‒, explaining why the SeCN‒/(SeCN)3

‒ and (SeCN)3‒/(SeCN)2

processes overlap significantly (see Figure 6, bottom). The decrease in E0’(X‒/X2) is offset by

33

the decrease in Kstab, and as a result E0’(X‒/X3‒) is 0.03 V more positive for the SeCN‒/(SeCN)3

‒

couple compared to the I‒/I3‒ couple.

From a thermodynamic perspective, the Br‒/Br3‒ couple and to a lesser extent SeCN‒

/(SeCN)3‒ couple are advantageous over the I‒/I3

‒ couple as a redox shuttle system in DSSC

technology, as they possess more positive formal (reversible) potentials. In practice, the open

circuit voltage of the DSSC is just one of many considerations which must be taken into account

when choosing a redox shuttle system. Other considerations include; (i) stability, corrosivity,

volatility and toxicity; (ii) mass-transport (high diffusivity minimizes charge-transport

limitations in the electrolyte) and; (iii) recombination kinetics with the electrons in TiO2. The

last point is particularly pertinent, as it is the very slow recombination kinetics between I3‒ and

electrons in TiO2 that sets I‒/I3‒ apart from most alternative (e.g., more strongly oxidizing)

redox mediators.

34

Conclusions

The formal potential of I‒/I2 redox couple and stability constant of I3‒ has been

quantified in two protic solvents, two aprotic solvents, eight aprotic ionic liquids and one protic

ionic liquid using a voltammetric method. E0’(I‒/I2) and Kstab(I3‒) were found to increase in the

order ethanol ≈ acetonitrile < propylene carbonate < AILs < PILs < water; and water < ethanol

≈ PILs < acetonitrile ≈ AILs < propylene carbonate, respectively. In the AILs, E0’(I‒/I2) and

Kstab lie in a narrow range (‒0.06 to 0.02 V vs. Fc/Fc+ and 106.1 to 107.1, respectively), while in

the PIL, [NH,2,2,2][NTf2], E0’(I‒/I2) is significantly more positive (0.11 V vs. Fc/Fc+) and Kstab is

smaller by approximately two orders-of-magnitude (104.7). Using the termolecular mechanism

proposed in our previous work and the values for DI‒, DI3‒, DI2, E0’(I‒/I2) and Kstab derived

voltammetrically here, the I‒ oxidation process was successfully modelled in all of the

investigated ILs. In [C2mim][NTf2], the basic DSSC additive, t-BPy, was found to shift E0’(I‒

/I2) negatively and decrease Kstab by two orders-of-magnitude, while another common DSSC

additive, Li+, was found to have a negligible impact on these parameters. Finally, in

[C2mim][NTf2], E0’(X‒/X2) and Kstab(X3

‒) were found to increase in the order SeCN‒/(SeCN)2

≈ I‒/I2 < Br‒/Br2 and (SeCN)3‒ << Br3

‒ < I3‒, respectively. The trends in

E0’(X‒/X2) and Kstab(X3‒) were rationalized in terms of the physicochemical parameters (i.e.,

polarity, Gutmann donor/acceptor numbers, ionic strength etc.) of the respective solvents/ILs.

In summary, we have shown that the parameter, E0’(X‒/X3‒), which directly influences the

operating voltage generated by a DSSC under illumination, can be tuned to an extent through

careful selection of the solvent and/or IL (i.e., constituent anion/cation). Although there are

many other factors (in addition to the energetics of the X‒/X3‒ process) which much be

considered when choosing the ideal combination of solvent/redox mediator in a DSSC

electrolyte, the work presented here will undoubtedly provide access to more informed

decisions on the matter.

35

Associated Content

Supporting information. Experiment-simulation comparisons of the iodide oxidation

process in aqueous media (Figure S1) and various ILs (Figures S2 to S8); and double-step

chronoamperograms used to estimate the parameters DBr3‒, DBr2

, D(SeCN)3‒ and D(SeCN)2

(Figures

S9 to S12). This material is available free of charge via the Internet at http://pubs.acs.org.

Author Information

Corresponding Authors

*E-mail: tony.hollenkamp@csiro.au (A.F.H) and pmahon@swin.edu.au (P.J.M)

Notes

The authors declare no competing financial interest.

Acknowledgements

C.L.B. acknowledges the financial support received from the Monash University

Faculty of Science Postgraduate Publication Award.

36

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44

FOR TOC ONLY:

Aqueous

Ethanol

Protic IL

Aprotic IL

Acetonitrile

E2 ‒

E1

Dec

reas

e Kstab

Increase

E (V)

E1

E2