Unveiling iodine-based electrolytes chemistry in aqueous dye-sensitized solar cells

The chemistry behind the I–/I3– redox couple is thoroughly investigated in 100% aqueous dye-sensitized solar cells, paving the way to this emerging green PV technology.

. Intensities of I 3 − signals of diluted aqueous solutions obtained from the aqueous electrolytes containing different amounts of NaI or KI salts (in M = mol L −1 ) at different times after the preparation of stock solutions. Amount of I 2 was fixed at 50 mM.

Stock solution
Abs ( Table S4. Equilibria involving aqueous solutions containing I 2 and I − .     Table S7. Parameters obtained by fitting the impedance spectra of aqueous DSSCs with the constraint:  Table S8. Parameters obtained by fitting the impedance spectra of aqueous DSSCs with the constraint:   Figure S2. Specific conductance of aqueous NaI and KI solutions resulting from the equivalent conductance data at different Iconcentrations extracted from reference [2]. More recent data (from reference [3]) referred to aqueous KI solutions are shown for comparison.   mM. Devices #2 and #3 were sealed after drying the entry slit to avoid any possible contamination of the electrolyte by glue. Lines just connect data points and have no physical meaning.

Appendix 1 -Determination of iodine species in aqueous electrolytes
The concentrations of I 3and Iare usually calculated from Eq. 2 considering the equilibrium in Eq. 1: where K = 723 L mol -1 in H 2 O and K = 5.8 × 10 6 L mol -1 in acetonitrile [4,5]. Therefore, in this work the equilibrium concentrations of I 3 -, Iand I 2 were calculated according to Eq. 2 assuming that the activity equilibrium constant equals K = 723 L mol -1 at 25 °C and that the activity coefficients ratio is equal to one. The literature value of K we used was determined spectrophotometrically at low Iand I 2 concentrations [6]. Palmer et al. found that the values of I 2 activity coefficient ( ) and the ratio ( ) vary by about the same quantities at should be equal to ≈1 over the investigated range [7].
Equilibrium concentrations of I 3and Iobtained as described above were used to calculate the redox potential of the electrolytes according to the Nernst equation: (Eq. 3) [ -] 3 assuming = 0.536 V [4]. In this case, the use of the standard potential instead of ± -3 ± - [7], errors up to 10 mV may be expected. The activity coefficients of Iin highly concentrated solutions would be required for accurate estimations. Pitzer parameters (e.g., parameters useful to understand the behavior of ions dissolved in water [8]) are compiled for NaI and KI solutions, but the expression returning the mean activity coefficients requires the value of ionic strength, which is affected by the formation of ion pairs. The evaluation of the activity coefficient of I 2 in salt solutions is often extrapolated from solubility data [9]. Such data are available in a wide range of Iconcentrations, but the formation of complexes does not allow to use the salting-out approximation.
For electrolytes containing I -0.50 M and I 2 50 mM, the equilibrium concentrations of higher polyiodides and the species following from the formation of hypoiodous acid (HIO) were computed on the basis of the algorithm developed by Gottardi [5]. The activity coefficient of uncharged species was set to equal one, and the activity coefficient of charged species was calculated by Debye-Hückel equation with an additional empirical term. Such assumptions allowed the author to obtain good agreement between the experimentally determined concentrations and the calculated ones (± ≈2% for potentiometric/titrimetric determinations and ± ≈5% for photometric measurements) up to I -≈0.60 M [10].
The equilibrium concentrations of iodine species were computed as a function of the pH.
At Iconcentrations higher than 0.50 M, the Debye-Hückel approximation is not suitable to calculate the activity coefficients of the charged species and sizeable errors are expected assuming the activity coefficient of uncharged species such as HIO and I 2 equals to one.
The activity coefficients were computed as follows: (Eq. 14) The five unknowns [I 2 ], [I -], μ, fa and ft were computed at each of the given pH values solving a system including Eqs. 16-18 and the two formulas that return the activity coefficients from Eqs. 14-15. All these equations were set to equal zero and the solutions were computed iteratively between bounds 0 ≤ x 1 ≤ C, 0 ≤ x 2 ≤ T, 0 ≤ x 3 ≤ T + 7C, 0 ≤ fa ≤ 2, 0 ≤ ft ≤ 2. All of the other concentrations were computed according to the previously listed equilibria.
The absolute value of the residual, |Δ|, was calculated as the difference between the sum of iodine atoms in all of the computed iodine species and T (