The electrolyte solutions were prepared in molality m (mol kg−1) concen- tration units. Molality was preferred over concentration c in mol dm−3 be-
cause the solution density changes with the salt present (see Fig. 99 in Ap- pendixG), and this effect is relatively more significant at lower salt concen-
5.2 preparation of the salt solutions 91
trations. As the preparation of solutions by molality involves weighing both solute and solvent, it does not present the above mentioned complication.
A kern® ABJ−220−4M analytical balance was therefore used. A Petri
dish was employed as support for the vial and to distribute its weight on a larger surface of the balance plate. The vials where the solutions were going to be prepared were washed with Type I ultrapure water, dried in an oven overnight and kept in a plastic desiccator over silica gel and under vacuum before use. Typically20 gto25 gof solution (depending on its density) were prepared each time in a 25 mlvial. This quantity was preferred over larger quantities in order to facilitate storage in desiccators. Also, to avoid prob- lems associated with storage (e.g. contamination, evaporation of the solvent), it was preferred to re-prepare the solution more often rather than storing it for a long time.
The salt was first weighed directly into the vial, followed by the addition of the solvent (as close as possible to the desired weight). Both salt and solvent weight were noted in order to calculate the solution concentration with the maximum accuracy achievable. The weighing procedure was performed in air, as quickly as possible. The usage of a glove bag was considered and trialled but the problems associated with the scale instability overcame the advantage of weighing in a dry atmosphere. The relative humidity in the laboratory was on average low (around25%), therefore swift operation and the monitoring of water content through Karl Fischer (kf) titration were preferred for practicality.
Parafilm M® was applied around the vial cap immediately after prepar- ation to help prevent moisture ingress. When dissolution was not instant- aneous, a vortex mixer and sonication were used to help the dissolution process. In some cases, such as for iodide salts, sonication was applied only briefly to avoid oxidation of the iodide to iodine (revealed by the appearance of a yellow colour). This is particularly true for dmso solutions of NaI: no sonication was applied at all in this case.
5.2.1 Solubility of the salts
The solubility of the electrolytes in the different solvents is listed in Table24. Most of the information was available in the literature. A salt concentra- tion that was accessible for most electrolytes in the different solvents is 0.05 mol kg−1. Therefore this was set as the electrolyte concentration for the
chromatography experiments. Qcm measurements were performed with electrolyte solutions of10−3mol kg−1 concentration. This lower concentra- tion was sufficient to produce detectable changes in the polymer brush be- haviour, whilst higher concentrations had little effect on the magnitude of the changes observed (see Fig.29in Chapter7).
Table 24: Literature solubilitySdata for the electrolytes investigated.
methanol
electrolyte S/mol kg−1 t/◦C reference
NaOAc 1.95* 15 H. Stephen and T. Stephen,1963
NaF 0.099* 20 H. Stephen and T. Stephen,1963
NaCl 0.24† 25 Pinho and Macedo,2005
NaBr 1.63† 25 Pinho and Macedo,2005
NaI 5.20* 25 H. Stephen and T. Stephen,1963
NaClO4 4.194 25 Chan et al.,1996 NaSCN 4.939 24.7 Hála,2004
formamide
electrolyte S/mol kg−1 t/◦C reference NaOAc >0.05k
NaF 0.026‡ 25 Scrosati and Vincent,1980
NaCl 1.61§ 25 Scrosati and Vincent,1980
NaBr 3.43§ 25 Scrosati and Vincent,1980
NaI 4.00§ 25 Scrosati and Vincent,1980
NaClO4 >0.05k
NaSCN 17.7§ 25 Scrosati and Vincent,1980
dimethyl sulfoxide
electrolyte S/mol kg−1 t/◦C reference
NaOAc 0.0078‡ 25 J. N. Butler,1967 NaF insoluble 25 J. N. Butler,1967
NaCl 0.08 25 J. N. Butler,1967 NaBr 0.55 25 J. N. Butler,1967 NaI 1.0 25 J. N. Butler,1967 NaClO4 1.8 25 J. N. Butler,1967 NaSCN 0.12‡ 25 J. N. Butler,1967 propylene carbonate
electrolyte S/mol kg−1 t/◦C reference NaOAc <0.05k NaF 5×10−5 25 Harris,1958 NaCl 3×10−6 25 Harris,1958 1.7×10−4‡ 25 Muhuri et al.,1993 NaBr 0.08 25 Harris,1958 3.6×10−3‡ 25 Muhuri et al.,1993 NaI 1.11 25 Harris,1958
NaClO4 2.5‡ 25 Muhuri et al.,1993 NaSCN >0.05k
*Calculated by the author, original given as ‘percentage by weight’Wt.%. †Calculated by the author, original given as mass fractionw
salt.
‡Units:mol dm−3.
§Calculated by the author, original given asg kg−1 solvent.
kno literature data available, the value listed comes from experience and refers to room temperature.
5.3 the influence of trace quantities of water 93
Table 25:Summary of the electrolytes investigated per solvent.
electrolyte meoh fa dmso pc NaOAc X X NaF X NaCl X X X NaBr X X X NaI X X X X NaClO4 X X X X NaSCN X X X X
The electrolytes investigated for each solvent are summarised in Table25. Although there are literature reports (Harris, 1958) that NaBr is soluble in pc at a concentration of 5×10−2mol kg−1, it was not possible to achieve its solubilisation at that concentration nor at the lesser concentration of 10−3mol kg−1. Note the solubility of salts varies greatly across solvents and is in most cases lower than in water. This limits the range of electrolytes and the concentrations available for investigation.
5.2.2 Storage of solutions
The iodine and thiocyanate solutions vials were covered in aluminium foil to prevent light degradation of the anions. All the vials containing solutions were kept in a glass desiccator over silica gel and granular phosphorous pentoxide, P4O10, which were changed regularly. Additionally, the desic-
cator was placed under a slight vacuum.