J. Phys. Chem. 1994,98, 6195-6200
6195
Spectroscopic Determination of Flatband Potentials for Polycrystalline Ti02 Electrodes in Mixed Solvent Systems Brendan Enright, Careth Redmond, and Donald Fitzmaurice' Department of Chemistry, University College Dublin, Dublin 4, Ireland Received: February 24, 1994; In Final Form: April 17, 1994'
We report studies directed toward a quantitative understanding of the relationship between the flatband potential (V,) of a metal oxide semiconductor electrode (titaniumdioxide) and the composition of a contiguous electrolyte solution (water, methanol, ethanol, and acetonitrile each containing 0.2 mol dm-3 tetrabutylammonium perchlorate). A relationship between Vhand the proton-generating ability of the electrolyte solvent, as measured by the autoprotolysis constant of the pure solvent, is demonstrated. Additional studies have examined the dependence of Vfion the composition of a binary electrolyte solvent mixture (acetonitrile/ethylene carbonate and acetonitrile/water). We conclude that the composition of the acetonitrile/ethylene carbonate mixture at the semiconductor-electrolyte solution interface (SLI) is similar to that of the bulk solution. However, for an acetonitrile/water mixture there is preferential adsorption of water at the SLI. Finally, for aprotic electrolyte solutions, the dependence of Vhon the electrolytic solute was studied. Cations such as Li+ and Na+ are potential determining.
Introduction Solar energy storage devices based on polycrystalline semiconductor electrodes are potentially important. However, practical devices are likely to require use of nonaqueous electrolyte solventsand visible light-sensitized wide bandgap polycrystalline semiconductor (metal oxide) electrodes. Nonaqueous solvents are preferred as they offer a wider range of electrochemical stability and greater solubility for electrolyticsolutes. Sensitized wide bandgap polycrystalline semiconductor electrodes are preferred as they are stable and cost effective. Recently, a highly efficient regenerative photoelectrochemical cell, based on visible light sensitizedpolycrystallineTi02 electrodes,has been described by O'Regan and Graetzel.1 Further development of such cells will require detailed understanding of how the band energetics of a polycrystalline metal oxide semiconductor electrode depend on the composition of a nonaqueous electrolyte solution. This, in turn, will necessitate determination of the flatband potential (Vh) of such electrodes under relevant operating conditions. Only limited data of this type are currently available for the following reasonx2 First, previously available methods for determining Vficould only be applied to single-crystal semiconductor electrodes, and second, Vh determinations in nonaqueous solvents are not easily reproduced. The first of these difficulties has been overcome following development of two spectroscopic methods which permit determination of Vh for transparent polycrystalline semiconductor ele~trodes.3.~The first of these methods measures, by visible spectroscopy, the number of electrons present in the conduction band at a given applied potential.' The difference between the applied potential and Vh, consistent with the spectroscopically determined concentration of conduction band electrons, is calculated by a fit to the experimental data. The second method may be applied if the density of conduction band states is such that an absorbance by free conduction band electrons is not mea~urable.~Under these conditions, absorbance changes at wavelengths shorter than those corresponding to the bandgap energy (Eg) may be measured. The observed absorbance loss, assigned to the Burstein shift accompanying band filling, is measured as a function of the applied potential. A graphical treatment is used to determine Vh. Abstract published in Advance ACS Abstracts, May 15, 1994.
0022-365419412098-6195%04.50/0
The second difficulty has largely been overcome following the realization that Vh in aqueous electrolyte solutions is determined by a proton adsorption-desorption equilibrium established at the metal oxide semiconductor electrodeliquid electrolyte solution interface (SLI).2 Absence of such an equilibrium in aprotic electrolyte solutions results in V, being highly sensitive to trace amounts of water and other protic impurities. Consequently, reproduciblevaluesfor Vhare only obtained in aprotic electrolyte solutions by using solutes and solvents which have been carefully dried and purified. An example of work of this type is determination of Vh for Ti02 in MeCN by Schumacher and co-w0rkers.~*6For very dry, highly pure, MeCN these workers observed a value for Vh of -2.2 f 0.1 V (SCE). This value is significantly more negative than that determined for Ti02 in MeCN containing 2% by vol added H20, i.e., -1.35 V (SCE). This shift of Vh to more positive values was attributed to trapping of positive charge at the SLI. In a recent study, we have used the first of the two spectroscopic methods outlined above to determine Vh for transparent polycrystalline Ti02 electrodes in a range of electrolyte solutions prepared from tetrabutylammonium perchlorate (0.2 mol dm-3) dissolved in dry acetonitrile (MeCN), ethanol (EtOH), and methanol (MeOH)? For water and nonaqueous protic solvents (MeOH and EtOH), consistent with establishment of a proton adsorption-desorption equilibrium in protic solvents at the SLI, Vh is significantly more positive than for nonaqueous protic solvents (MeCN).2 It was also observed that for water and nonaqueous protic solvents Vhis independentof chosen electrolyte, whereas in nonaqueous aprotic solvents V, depends on the electrolytic solute. Cations which were found to be potential determining were Li+, Na+, and Mg2+, Some evidence was presented which suggested that at low cation concentration adsorption is responsible for the observed positive shift in V,, while at higher cation concentrations intercalation close to the electrode surface may be important. Here we report work directed toward a more quantitative understanding of the relationship between the proton-generating ability of the electrolyte solvent and Vh. The dependence of Vfi on the composition of mixed electrolyte solvent systems is also studied in an attempt to further extend our understanding of those factors controlling Vh. Finally, the mechanism by which cations such as Li+, Na+, and Mg2+ are potential determining in single and mixed solvent systems is examined. 0 1994 American Chemical Society
6196 The Journal of Physical Chemistry, Vol. 98, No. 24, 1994
Enright et al.
a
TABLE 1: Flatband Potentials of Polycrystalline Ti02 Electrodes in Electrolyte Solvent Containing 0.2 M Tetrabutylammonium Perchlorate and Added Lithium Perchlorate 0.0
MeCN EtOH MeOH
-2.04 -1.39 -1.12 HzO (PH1l)d -1.06 H20 (PH 2) -0.52 EC -1.59
O-r----7 n
W
[LiC104],M 1 X l e 3 1 X 10-1 pK'(25 0C)6 pK,(molar)C V;b -1.97 -0.90 26.5 29.49 -1.24 -1.37 -1.10 19.1 21.98 -0.99 -1.11 -1.10 16.7 19.90 -0.95 -1.07 14.W 17.49e -0.82e -0.49
W " ,
> e >
v
Slope = 9 7 n b i l p K s
10
All flatband potentials are in volts versus SCE.* Values taken from ref8. e Valuescalculatedaccordingtoeqs12and 13. Aqueouselectrolyte solutions were prepared with 0.1 M tetramethylammoniumperchlorate. e At pH 7.0.
Experimental Section Measurement of Flatband Potentials. All electrodes consisted of a 4-pm-thick layer of fused Ti02 particles (12-nm diameter) supported on a 0.5-pm-thick layer of fluorine-doped SnOz on glass. Preparation of theseelectrodes has been described in detail elsewhere.' Briefly, Ti02 was prepared by hydrolysis of titanium isopropoxide. The resulting dispersion was concentrated to about 160 g/L and Carbowax 20000 (40 wt % equivalent of TiO2) added yielding a white viscous liquid used to form a 4-pm-thick layer on conducting glass. After drying in air for 1 h, each electrode was fired a t 500 OC for 12 h and stored in a vacuum desiccator. A Ti02 film formed the working electrode (2-cm2surface area) of a closed three-electrode single compartment cell, the counter electrode being platinum and the reference electrode a saturated calomel electrode (SCE) connected via an appropriate salt bridge. Potential control was provided by a Thompson Electrochem Ministat potentiostat connected to a HP 33 10A function generator. The above cell was incorporated into the samplecompartment of a Hewlett Packard 8452A diode array spectrometer. Absorbance was measured a t 356 and 780 nm, the applied potential being scanned at 5 mV s-l. For each determination of V , a new working electrode and freshly prepared electrolyte solution were used. Measurement of Electrophoretic Mobilities. Electrophoretic mobilities were determined a t 25.0 f 0.1 OC using a Rank Bros. Mark I1 instrument (thin-walled cylindrical cell configuration). Particle velocity was determined by following the motion of a Ti02 particle over 114 pm. The electrophoretic velocity was divided by the field strength to yield the electrophoretic mobility p. The mobility was determined 20 times within each stationary layer, Le., 10 times in each direction. The mean value for p was used in subsequent calculations. Samples were prepared by sonication of Ti02 particles, having an average diameter of 12 nm, into a volume of dry solvent to yield a dispersion containing 0.01 g/L of TiO2. (To ensure these particles were similar to those constituting the semiconductor electrodes, they were fired at 400 O C for 12 h prior to dispersion by sonication.) Electrolytic solute was added to a concentration of 10-3 mol dm-3. Measurement of Adsorption Isotherms. Polycrystalline films (4 pm thick and 12.4 mg TiOZ/slide) were prepared on glass microscope slides (geometric surface area of 17.7 cmz) by firing at 500 OC for 12 h. Five films were placed in a staining bath and 120 mL of dry solvent containing a known initial concentration (CO)of electrolytic solute. After 24 h the solution was sampled for analysis by atomic absorption spectroscopy, and the concentration of electrolytic solute remaining in the solution phase (C,) was determined. In each case a simultaneous blank experiment utilizing five uncoated microscope slides was performed. In no case, for which data are reported, was significant adsorption detected in the blank experiment. heparation of Solvents and Solutes. Aqueous electrolyte solutions were prepared using distilled deionized water (pH
1
-3 7 15
30
25
20
PKs
b n
w
O
c",
-1
W
>
r
i
MeOH
v
-e
MeCN
>
I I
-2 15
Slope = 34 mV/pKs(molar)
20
25
30
pK, ( molar) Figure 1. (a) Flatband potential determined for a polycrystalline Ti02 electrode in electrolyte solution containing 0.2 mol dm-3 TBAP (0.1 M TMAP in HzO)plotted against pK, of electrolytesolvent. (b) Flatband potential (normalized against molarity of solvent) plotted against pK,(molar) for data in (a). (Data are taken from Table 1.)
adjusted used KOH or HClO,). MeCN and ethylene carbonate (EC) were refluxed over P205 for 3 hand recovered by distillation, and themiddle fractions were retained for immediate use. EtOH was refluxed over magnesium ethoxide for 3 h and recovered by distillation, and the middle fraction was retained for immediate use. MeOH was refluxed over CaH2 for 3 h and recovered by distillation, and the middle fraction was retained for immediate use. Tetrabutylammonium perchlorate (TBAP), tetramethylammonium perchlorate (TMAP), lithium perchlorate, and sodium perchlorate were used following drying under vacuum a t 150 OC for 48 h. Rt?SUIts
Flatband Potentials. Generally, Vb for a Ti02 electrode in aqueous electrolyte solution depends on pH, with Vb being given by eq 1.3
V, = -0.40 - 0.060pH [V,SCE] (1) By plotting absorbance changes at 780 nm, assigned to accumulation of conduction band electrons, on the same potential scale as those measured for an aqueous electrolytic solution of known pH, it is possible to determine V b for a nonaqueous electrolytes ~ l u t i o n .The ~ assumptions underlying this treatment are discussed below. Specifically, V b has been determined for MeCN, EtOH, and MeOH (0.2 mol dm-3TBAP).7 These results, and those obtained for water (0.1 mol dm-3 TMAP), are summarized in Table 1. Also in Table 1 are values for Vb determined following addition of 10-1 and 10-3mol dm-3 LiClOd and literature values for the autoprotolysis constants (KJ of the above solvents.* A plot of V b against pK, yields the straight line relationshipshown in Figure la. To examine the effect of solvent composition, Vb has been determined as a function of the mole fraction of MeCN for the following mixed solvent systems: MeCN/HzO (0.2 mol dm-3 TBAP, except in pure H20 with 0.1 mol dm-3 TMAP) and MeCN/EC (0.2 mol dm-3 TBAP). The results of these experiments are summarized in Tables 2 and 3
The Journal of Physical Chemistry, Vol. 98, No. 24, 1994 6197
Flatband Potentials for Polycrystalline Ti02 Electrodes
a
TABLE 2 Flatband Potentials of PolycrystallineTiOt Electrodes in a Mixed Acetonitrile/Water Electrolyte Solvent Containing 0.2 M Tetrabutylammonium Perchlorate and Added Lithium Perchlorate
n
mole fraction MeCN
0.0
1.oo 0.97 0.75 0.50 0.25
-2.04" -1.81 -1.39 -1.24 -1.12 -0.82
0.00 (pH 7)d
(-2.04)b (-2.03) (-1.74) (-1.43) (-1.13) (-0.82)
%
1x10-3 I X ~ O - - L (-2.04)' (-1.81) (-1.41) (-1.20) (-1.11) (-0.82)
-1.97
-0.90
-1.39 -1.24 -1.12 -0.82
-1.02 -0.92 -0.84 4-82
LI = 0.001 M
W
ILiClOd. M
>-
v
........ -0......__._
e
>
0
"All flatband potentials are in volts versus SCE. Mole fraction weighted sum of flatband potentials determined for pure solvents; see Table 1. Surface coverage weighted sum of flatband potentials determined for pure solvents;see Table 1. d Aqueous electrolyte solutionswere prepared with 0.1 M tetramethylammonium perchlorate.
1 X(MeCN)
b - 3 LI
-
+
0.000M
TABLE 3: Flatband Potentials of Polycrystalline Ti02 Electrodes in a Mixed Acetonitrile/Ethylene Carbonate Electrolyte Solvent Containing 0.2 M Tetrabutylammonium Perchlorate and Added Lithium Perchlorate [LiC104], M mole fraction MeCN
0.0
1x10-3
1x10-1
1.oo 0.75 0.50 0.25
-2.04" (-2.04)b -1.92 (-1.93) -1.83 (-1.81) -1 -72 (-1.70) -1.59 (-1.59)
-1.97 -1.86 -1.77 -1.66 -1.52
-0.90 -0.88 -0.88 -0.75 -0.76
0.0w
-2w
0 All flatband potentials are in volts versus SCE. Mole fraction weighted sum of flatband potentials determined for pure solvents; see Table 1. Electrolyte solution maintained at 50 OC.
-3
MeCN/EC
n
w 0 c",
>
W
e
>
"
I
0
1 X(MeCN)
Figure 3. Flatband potential determined for a polycrystalline Ti02 electrode in electrolyte solution containing 0.2 mol dm-3 TBAP (0.1 M TMAP in pure H20) plotted against mole fraction of MeCN in mixed solventsystem (a) MeCN/H20 and (b) MeCN/EC. Values determined mol dm-3 for the flatband potential following addition of 10-l and LiClO4 are also plotted. (Data are taken from Tables 2 and 3.)
0
TABLE 4
- 1 -----*----- 0 0
1
x (MeCN) Figure 2. Flatband potential determined for a polycrystalline Ti02 electrode in electrolyte solution containing 0.2 mol dm-3 TBAP (0.1 M TMAP in pure H20) plotted against mole fraction of MeCN in mixed solventsystem. The solid line represents the predicted flatband potential for the MeCN/EC system based on a mole fraction weighted sum of the flatband potentialsofthe puresolvents (eq 12). Thedashedlinerepresents the predicted flatband potential for the MeCN/H20 system based on a surface coverage (Langmuir isotherm) weighted sum of the flatband potentials of the pure solvents (eq 14). (Data are taken from Tables 2 and 3.)
and plotted in Figure 2. The effect on Vfiof addition of and 10-1mol dm-3 LiC104, as a function of the mole fraction of MeCN in each of the above solvent systems, has also been examined. The results of these experiments are summarized in Tables 2 and 3 and plotted in Figure 3. Electrophoretic Mobilities. Electrophoretic mobilities have been determined for Ti02 nanocrystallites in MeCN, EtOH, MeOH, and HzO. p was also determined, in each of the above solvents,following addition of LiC104,NaC104,and TBAP ( mol dm-3); see Table 4. The corresponding zeta potentials (0 were calculated using eq 2.9
ANU)is determined by interpolation between tabulated values,
Electrophoretic Data for Ti02 Nanocrystallites
electrolyte
mobility [X lo8] p (M2 S-*V-l 1
blank LiC104 NaClO4 TBAC104
-1.84 -1.84 -1.83 -1.90
blank LiC104 NaClO4 TBAClO4
-1.89 -2.01 -1.92 -2.72
blank LiC104 NaC104 T BA C104
-0.82 -0.55 -0.63 -0.59
blank LiCIO4 NaCIO4 TBAC104
-3.60 -1.82 -2.36 -3.71
zeta potential C (mv)
charge [X 1015] u (e M-2)
Solvent: Water -35 -3 5 -3 5 -36
2.01 3.25 3.25 3.34
Solvent: MeOH -54 -56 -53 -75
1.29 2.57 2.45 3.48
Solvent: EtOH -62 -41 -46 -43
1.10 1.49 1.68 1.60
Solvent: MeCN -58 -29 -38 -58
1.52 1.42 1.89 2.97
Averageparticlediameter 12 nm. Concentrationof added electrolyte mol dm-3. Concentration of added Ti02 equal to 0.01 g equal to dm-3.
where K is the Debye length, a is the particle radius, and 9 is the solvent viscosity.10 For the pure solvents, containing no added electrolyte, the ionic strength is sufficiently small thatAKa) may be assumed to be zero. For solvents containing added electrolyte AKU)is calculated by taking into account the degree of dissociation of each electrolyte in a given solvent. The total particle charge (Q)is calculated using eq 3 (see Table 4).
Enright et al.
6198 The Journal of Physical Chemistry, Vol. 98, No. 24, 1994
Q = {ea(l
+ KU)
(3)
The correspondingsurface charge density (a) is calculated using eq 4 and expressed in terms of the number of electronic charges per unit area; see Table 4. u
= Q/4aa2
(4)
Adsorption Isotherms. Langmuir adsorption isotherms have been prepared according to eq 5, using the data summarized in
Table 5 for adsorption of LiC104 from MeCN, EtOH, MeOH, and H20. n is C, - C,, C, is the adsorbate concentration at monolayer coverage, and K is an equilibrium constant given by the ratio of the rate constants for adsorption and desorption(kah/ kds). The above isothermsare shown on different scales in Figure 4. It is clear Li+ ions are adsorbed from MeCN and EtOH. No significant adsorption of Li+ ions is observed from MeOH or H20. Similar experiments with NaC104 showed no significant adsorption of Na+ in MeCN, EtOH, MeOH, and H20. Discussion Dependence of V, on the plv, of the Electrolyte Solvent. Vfi for a Ti02 electrode in contact with aqueous electrolye solution depends on the proton adsorption4esorption equilibrium estab lished at the SLI.2 Vfi determinations for polycrystalline Ti02 electrodes, in contact with aprotic electrolyte solvents, are therefore sensitive to trace amounts of water and reported values have shown significant variation.2 To better understand the dependenceof Vbon the properties of the electrolyte solvent, this quantity has been determined for a range of carefully dried and purified organic solvents and water;2.5.6 see Table 1. These results are discussed in detail below. The Galvani potential difference between the bulk of the semiconductor electrode and the electrolyte solution (A+) is determined by the potential drop across the space charge layer of the semiconductorelectrode (Qsc), the Helmholtz layer (@H), and the space charge region of the electrolyte solution (*/EL), eq 6-11
A 4 = q s c + 9, + *EL
TABLE 5 Adsorption Data for Lithium Perchlorate at P o l ~ c ~ s t a l l i nTi02 e Films Co [mol dm-31 C, [mol dm-31 n [mol dm-31 C,/n Solvent: Water 1.00x 10-4 1.00 x 10-4 0.00 2.50 X 5.00 X 7.50 X 1.00 x 1.00 x 5.00 x 1.00 x
10-4 10-4 10-4 10-3 10-2 10-2 10-1
2.50 X 5.00 X 7.50 X 1.00 x 1.00 x 5.00 X 1.00 x
1.00 x 2.50 X 5.00 x 1.00 x 1.00 x 5.00 x 1.00 x
10-4 10-4 10-4 10-3 10-2 10-2 10-1
1.00 x 2.50 X 5.00 X LOO x 1.00 x 5.00 x 5.00 x
1.00 x 2.50 X 5.00 X 1.00 x 1.00 x 5.00 X 1.00 x
10-4 10-4 10-4 10-3 10-2
1.00 x 2.50 X 5.00 x 1.00 x 1.00 x 4.87 X 0.98 X
1.00 x 2.50 X 5.00 X 7.50 X 1.00 x 1.25 X 1.50 X 1.75 X 2.00 x 2.25 X 2.50 X 2.50 X 5.00 X 2.50 X 5.00 x 1.00 x
10-4 10-4 10-4 10-4 10-3
10-4 10-4 10-4 10-3 10-2 1CF2
lo-'
0.00 0.00 0.00 0.00 0.00 0.00 0.00
Solvent: MeOH 10-4 10-4 10-4 10-3 10-2 10-2 10-2
0.00 0.00 0.00 0.00 0.00 0.00 0.00
Solvent: EtOH
10-1
10-4 10-4 10-4 10-3 10-2 10-2 10-L
0.00 0.00 0.00 0.00 0.00 0.13 X 1CF2 0.02 x 10-1
37.5 57.8
Solvent: MeCN
le3 lk3 10-3
le3 10-3
le3 lP3 10-2 10-1
0.10 x 1.08 X 3.40 X 5.92 X 0.85 X 1.11 x 1.34 X 1.59 X 1.86 X 2.15 X 2.36 X 2.38 X 4.80 X 2.46 X 4.93 x 0.99 X
10-4 10-4 10-4 10-4 10-3 10-3 lk3 lW3
le3
lk3 lk3 10-3 10-2 10-2 10-1
0.90 X 1.42 X 1.60 X 1.58 X 0.15 X 0.14 X 0.16 X 0.16 X 0.14 X 0.10 x 0.14 X 0.12 x 0.20 x 0.04 X 0.07 X 0.01 x
10-4 10-4 10-4 10-4 lt3 lk3 lk3 lP3 lP3 10-3 lP3 10-3 10-3 1k2 10-1
0.1 0.8 2.1 3.8 5.6 8.1 8.6 10.1 13.4 20.4 16.6 20.4 24.1 58.7 68.4 115.0
as occurs upon lowering the pH in aqueous solution, will result in a change in *H and a shift in Vfito more positive values. More usefully, we may express \kH in terms of proton activity (eq 10).
(6)
+/EL may be neglected at the high electrolyte concentrations for which results are reported (0.2 mol dm-3), eq 7.
qH= constant For moderate degrees of accumulation, Le., small space-charge capacitances (CCS), and typical Helmholtz layer capacitances (CH)*H will be small, and values determined for Vh in different solvents may be compared (eq 8).
Generally, *H depends on the charge accumulated in the Helmholtz layer (qH) and on dipole orientation effects at the SLI (eq 9).12
Equation 9 predicts that specific adsorption of protons ( q H > O),
+ ( 4 0 5 9 ) p H + xsLI
(1Oc)
Equation 10 predicts V, will shift to more positive values by 59 mV/pH unit upon lowering pH. The observed pH dependence of V b for polycrystallineTi02 electrodes, given by eq 1,' agrees well with this prediction. On changing from water as the electrolyte solvent, it is reasonableto expect that *H and therefore Vfi will change. First, it is certain the effective pH of the solvent will differ from that of water. Second, the magnitude of any potential drop due to dipole orientation effects may change. A plot of Vfiagainst pK,, the latter being a measure of the proton-generatingability of the solvent, yields the linear relationship in Figure la. We find that Vfiis shifted to increasingly positive values as the proton-generating ability of the electrolyte solvent increases. This is strong evidence to support the view that adsorption of positive charge, specifically protons generated by the solvent, determines +H and therefore Vh. The slope of
The Journal of Physical Chemistry, Vol. 98, No. 24, 1994 6199
Flatband Potentials for Polycrystalline Ti02 Electrodes
a
30r---l u
component solvents is not significantly different from that of the respective pure solvents over the entire composition range. Using eq 13 to predict V , for MeCN/H20 clearly does not account for the values measured; see Table 2. We therefore conclude that the composition of MeCN/H20 at the SLI differs significantly from that in the bulk. The measured values for V, suggest an excess of H 2 0at the SLI and were fitted to a Langmuir adsorption isotherm, eq 14.
EtOH/MeOH/HOH
n
O.OE+OO
2.5E-03
eM&N is the degree of coverage by MeCN as a function of the mole fraction of MeCN in the MeCN/HzO, and K is a dimensionless constant equal to 4.36. V, is calculated as a coverage weighted sumof the flatbandpotentialsof thecomponent solvents (eq 15).
Ceq (mol -1 dm-3 )
b l 2 O p F MeOH/HOH - - \
V,(MeCN/EC)
= 8M&N(-2 u"
OE+OO
1E-01
Ceq (mol-\ dm-3) Figure 4. Langmuir adsorption isotherm for polycrystalline Ti02 films in electrolyte solvents containing added LiClO4 at concentrationsup to (a) 2.5 X le3and (b) 1.0 X le1 mol dm-3. (Data are taken from Table 5.)
the best fit line in Figure l a is 97 mV/pK, with a correlation coefficient (r2)of 0.99. The value of the sfope is apparently not consistent with the observation that in water Vh is shifted to more negative values by 59 mV/pH, equivalent to a shift of 29.5 mV/pK,. However, to compare values for V, measured in solvents having different standard states, we must first convert K, to K,(molal) by dividing by the square of the solvent molality (eq ll)." &(molal) = KS/[SHl2
(11)
The molar autoprotolysis constant is given by eq 12 K,(molar) = &(molal) - 2 log(po)
(12)
where po is solvent density. Additionally, Vh must also be corrected for solvent molarity. The resulting values V h are referenced to water. Upon doing this, we obtain a trace whose slope 34 mV/pK,(molar) with a correlation coefficient ( r z ) of 0.99, in good agreement with the expected value; see Figure 1b. An implication of the above being that \ t ~and , therefore V ~ O is , largely determined by the proton-generatingability of the solvent. V, and Binary Electrolyte Solvent Systems. Vfi of a polycrystallineTi02 electrodedepends on the compositionof the binary electrolyte solvent (BES). We first consider the case where mixed solvent compositionat the SLI and in the bulk solvent are similar. We predict Vb will equal the mole fraction weighted sum of the flatband potentials of the component solvents. Specifically, for MeCN/EC Vm is given by eq 13 where X M ~and N XECare the ) Vh(MeCN/EC) = X M , C N ( - ~ . ~ ~XEc(-1.59)
[ v , SCE] (13)
mole fractionsof MeCN and EC, respectively;see Table 3. Values calculated for V,, using eq 13, are in excellent agreement with the measured values; see Table 3 and Figure 2. On the basis of the above, we conclude that the proton-generating ability of the
The calculated values plotted in Figure 2 are in excellent agreement with the measured values, provided the mole fraction of MeCN exceeds about 0.2. For pure H20 the Vh is more positive than that predicted based on the values observed for the organic-rich solvent mixture. There are interesting parallels between these observations and those concerning solvating propertiesof MeCN/H20 reported by Waghorne et al.14J5 These workers observe a sharp discontinuity in transfer enthalpies at mole fractions for MeCN greater than about 0.2. This, and related observations, have been accounted for by a change in solvent structure from one based on three-dimensionalhydrogen-bonded water to one of lower order based on the organic comp0nent.~~J5 Equivalent discontinuities have not been observed by these workers for nonaqueous mixed solvent systems, as in MeCN/EC. The more positive than predicted value for Vmin pure H20 is, therefore, attributed to a change in solvent structure at the SLI and a correspondingchange in \ k ~ .The following reasons for a change in \ k due ~ to a change in water structure have been considered: There is a change in the proton-generating ability of one or both of the component solvents; the degree of TBA+ solvation varies with composition; the effective dielectric constant at the SLI varies with composition. Further work is in progress in this area. V, and Electrolytic Solutes. To date we have found V, to be independent of the electrolyeused in protic solvents, for example H20 and MeOH. However, for aprotic solvents, such as MeCN, Vh may depend on the cation of the added electrolyte. It was suggested this is due to the presence and absence of a proton adsorption-desorption equilibrium in a potential-determining fashion at the SLI. A proton adsorption-desorption equilibrium being assumed to preclude cation adsorption/intercalationat the SLI. Specifically,LiC104added to MeCN (0.2 mol dm-3 TBAP) results in a shift of Vhto more positive potentials, consistent with Li+ being the potential-determining ion. At Li+ concentrations up to mol dm-3 there is a gradual shift of about 40 mV/ molar decade.' At higher Li+ concentrations there is a rapid shift of Vh to more positive values. It was suggested that adsorption is important at lower Li+ concentrations and intercalation close to the electrode surface is important at higher Li+ concentrations.7 The positive shift in Vh following addition of NaC104is smallerthan that observed followingadditionof Lie104 at lower concentrations, but very much less than that observed at higher concentrations, consistent with the larger diameter of the Na+ ion precluding intercalation.' The electrophoreticmobility of Ti02 nanocrystallitesdispersed in a range of solvents containing added electrolyte support the above interpretation; see Table 4. In aqueous dispersionsthere is little change in { following addition of LiC104, NaC104 or TBAP, consistent with the absence of significant ion adsorption
Enright et al.
6200 The Journal of Physical Chemistry, Vol. 98, No. 24, 1994 in this highly protic solvent due to the existence of a proton adsorption-desorption equilibrium at the SLI. In MeCN t increasessignificantly following addition of LiC104 and NaC104, consistent with cation adsorptionin aprotic solventsdue to absence of a proton adsorption-desorption equilibrium at the SLI. We note the increase in { following addition of LiC104 ( l e 3 mol dm-3) is about twice that observed following addition of NaCl04 ( 10-3mol dm-3), Addition of TBAP results in only a small change in {. In the case of MeOH and EtOH, protic and moderately aprotic solvents, respectively, the expected intermediate behavior is observed. An exception is addition of TBAP to MeOH which results, reproducibly, in an increase in mobility. These studies confirm that addition of Li+ and Na+ ions to aprotic solvents reduces the surface charge density of a Ti02 nanocrystallite by adsorption and/or intercalation. No such effect is observed for proticsolvents. These observationscorrelate well with the effects of Li+ and Na+ addition on Vh of polycrystallineTi02 electrodes in protic and aprotic electrolyte solution^.^ Studies were performed to better understand cation adsorption by Ti02 in aprotic solvents. Up to Li+ concentrations of mol dm-3 adsorption by a polycrystallineTi02 electrode in MeCN is described well by a Langmuir adsorption isotherm; see Figure 4a. No adsorption of Li+in EtOH, MeOH, and H2O was detected in this concentrationrange. These findings support our conclusion that at low concentrations, in aprotic solvents, Li+ ions are adsorbed at the SLI shifting Vh to more positive values. At higher concentrations a complex isotherm is observed for Li+ in MeCN; see Figure 4b. This is assigned to intercalation of Li+ at the SLI."5J7 Adsorption of Li+ in EtOH is apparent at these higher concentrations,consistentwith the intermediate aproticity of this solvent, and is well described by a Langmuir isotherm. No adsorption of Li+ in MeOH or HzO is observed at these higher concentrations. Finally, no adsorption of Na+ ions was detected at any concentration in the above solvents. This is broadly consistent with the larger diameter of Na+ precluding intercalation and the smaller shift in Vh measured for this cation.'* It appears to be generally accepted that strongly hydrated ions, such as Li+, are preferentially adsorbed at anatase in aqueous electrolyte solution.19 It is apparently equally well established that this process occurs with retention of the Li+ hydration shell intact. The proposed driving force for adsorption is localized hydrolysis of a coordinating water molecule that facilitates attachment of Li+ to nonionized surface OH (basic) groups. The results presented in this paper suggestthat preferential adsorption of this type does not result in a shift of Vhof polycrystallineTi02 electrodesto more positive values. Thus, we conclude adsorption of hydrated Li+ ions is in a manner that does not affect 9 ~We. note that in the adsorption studies reported it was not possible to detect significant reductions in the bulk concentration of Li+ ions in aqueous solution. However, under similar conditions, it proved possible to detect significant reductions in bulk Li+ concentrationsfor MeCN. This suggestsa qualitativelydifferent mechanism for Li+ adsorption in aprotic solvents. The fact that adsorption is accompaniedby significant changes in Vhindicates these ions are adsorbed in a manner that changes \ k ~ Specifically, . that they are adsorbed at the electrode surface. Further support for this view is the significant changes in p that accompanies Li+ adsorption. Unlike in aqueous solution, cation adsorption results in a change in uof the hydrodynamicentity. Therefore,it appears that in aprotic solvents the absence of a proton adsorptiondesorption equilibrium irreversible adsorption/intercalation of Li+ ions is preferred to the apparently weak association of the solvated ion at an electrode in aqueous solution. As the protongenerating ability of the solvent increases, the tendency for a Li+ ion to desolvate and become adsorbed/intercalated at the electrode surface is reduced. Such a view being consistent with the behavior of Li+ ions in EtOH, a Solvent Of intwmediate proticity.
The effect of addition of LiC104 on Vhin mixed solvent systems was examined (Figure 3). For MeCN/EC addition of 10-3 and 10-l mol dm-3 LiC104 resulted in a shift of Vb to more positive values by an average of 0.06 and 0.98 V, respectively, across the composition range. Both MeCN and EC are aprotic solvents, and as expected the effect of added LiC104 is independent of composition. As above, for MeCN/H20 additionof LiC104 shifts Vh to more positive potentials. However, the magnitude of the change in Vh is dependent on composition, consistent with this being a mixed solvent system containing a highly protic (H20) and aprotic solvent (MeCN). As ~ M C N decreases, Vm will increasingly be dominated by establishment of a proton adsorption-desorption equilibrium. Consequently, the magnitude of the shift in V, to more positive values, following addition of 10-3 or le1mol dm-3 LiC104, also decreases until for an aqueous electrolyte solution no change in Vh is measured; see Figure 3.
Conclusions A relationship between Vh and the proton-generating ability of the electrolyte solvent, as measured by the autoprotolysis constant of the pure solvent, has been demonstrated. Additional studies have examined the dependence of Vfion the composition of a binary solvent electrolyte solution. We conclude that the composition of MeCN/EC at the semiconductor-electrolyte solution interface (SLI) is similar to that of the bulk solution. However, for MeCN/H20 there is preferential adsorption of water at the SLI. Finally, for aprotic electrolyte solutions, the dependenceof Vh on the electrolytic solute was studied. Cations such as Li+ and Na+ are potential determining. Regarding behavior of these ions at the SLI, adsorption and intercalation are important for Li+; however, only adsorption is important for Na+.
Acknowledgment. This work was supported by HewlettPackard (Ireland) Ltd and by EOLAS (The Irish National Agency for Science and Technology.) The authors thank Dr. Waghorne for his helpful comments. References and Notes (1) ORegan, B.; Graetzel, M. Nature 1991, 353, 737. (2) Finklea, H. Semiconductor Electrodes; Elsevier: New York, 1988. (3) Rothenberger, G.; Graetzel, M.; Fitzmaurice, D. J . Phys. Chem. 1992, 96, 5983. (4) Redmond, G.; Burgess, C.; OKeeffe, A,; MacHale, C.; Fitzmaurice, D. J . Phys. Chem. 1993, 97, 11081. (5) Heinzel, A. B.;Teschner, D. M.; Schumacher, R. Ber. Bunsen-Ges. Phvs. Chem. 1981.85. 1117. (6) Schumacher, R.;Teschner, D.; Heinzel, A. B. Ber. Bunsen-Ges.Phys. Chem. 1982,86, 1153. (7) Redmond, G.; Fitzmaurice, D. J . Phvs. Chem. 1993.97. 1426. (8j Coetzee, J. F.; Padmanbhan, G. R. i Phys. Chem. 1962,' 66, 1708. (9) Hunter, R. J. Zeta Potential in Colloidal Science; Academic Press: London, 1981. (IO) Abramson, H. A.; Moyer, L. S.;Gorin, M. H. Electrophoresis of Proteins; Reinhold: New York, 1942. ( 1 1 ) Green. M. J. Chem. Phys. 1959,31, 200. (12) Bard, A. J.; Fan, F. F.; Gioda, A. S.;Nagasubramanin, G.; White, G. S.Faraday Discuss. Chem. Soc. 1980, 70, 19. (13) King, E. J. In Physical Chemistry of Organic Solvent Systems; Covington, A. K., Dickinson, T., Eds.; Plenum Press: New York, 1973; pp 391-3. (14) Waghome, E. W.Chem. Soc. Rev. 1993, 285. (IS) Carthy, G.; Feakins, D.; ODuinn, C.; Waghorne, E. J . Chem. Soc., Faradav Trans. 1991. 87. 2447. ( 1 6 j Tennakone,K.;Wickramanyake, S.W.;Samarasekara,P.; Fernando, c. A. phys* ''Iidi 1987i '04* K57. (17) Murray, D. J.; Healy, T. W.; Fuerstenau, D. W. In Advances in Chemistry; Gould, R. F., Ed.; American Chemical Society: Washington, D.C., 1968; Vol. 79, Chapter 7. (18) Ooi, K.; Miyai, Y.; Sakakihara, J. Langmuir 1991, 7 , 1167. (19) Augustyoski, J. Aspects of Phot&electrochemical and Surface Behaviour of Titanium (IV) Oxide in Structure and Bonding 69; SpringerVcrlag: Berlin, 1988.
