Article pubs.acs.org/jced
Densities and Molar Volumes of Aqueous Solutions of LiClO4 at Temperatures from 293 K to 343 K Bin Hu,† Lubomir Hnedkovsky,‡ Wu Li,† and Glenn Hefter*,‡ †
CAS Key Laboratory of Salt Lake Resources, Chemistry, Qinghai Institute of Salt Lakes, Chinese Academy of Sciences, Xining 810008, China ‡ Chemistry Department, Murdoch University, Murdoch, Western Australia 6150, Australia ABSTRACT: Densities of aqueous solutions of purified lithium perchlorate at solute concentrations ranging from approximately 0.05 to 5 mol·kg−1 have been measured by vibrating-tube densimetry at 5 K intervals over the temperature range 293 ≤ T/K ≤ 343 at atmospheric pressure. Apparent molar volumes of LiClO4(aq) derived from these data varied smoothly with solute concentration and temperature and could be fitted with an average precision of ±0.02 cm3·mol−1 using an extended Redlich−Rosenfeld−Meyer equation. The present results show that volumetric data listed in popular literature compilations are inaccurate. The effects of temperature on the partial molar volumes of LiClO4(aq) are nonlinear and indicate that the structure-breaking or -making character of this salt depends on concentration and temperature. Systematic differences observed between the volumes of lithium and sodium salts with a common anion appear to be related to differences in the strength of hydration of the two cations.
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INTRODUCTION Lithium perchlorate is widely used as a noncomplexing “swamping” electrolyte to control activity coefficients in equilibrium, kinetic, spectroscopic, and electrochemical studies in aqueous solutions.1,2 It is especially useful for this purpose in strongly acidic mixtures with perchloric acid,3 although its noncomplexing credentials are somewhat dubious (because Li+(aq) often binds anions to a significant extent).4 In biochemistry, saturated LiClO4(aq) solutions are used to denature proteins, while the high solubility of LiClO4 in many organic solvents has led to its use in lithium-ion batteries and as a catalyst in several named organic reactions. Regardless of the widespread use of LiClO4(aq) solutions, their volumetric properties are not well established. Early measurements on the densities (ρ) of LiClO4(aq) have been summarized by Timmermans.5 Thus, Mazzuchelli and Rossi6 reported densities of LiClO4(aq) at three concentrations and two temperatures (288 and 298 K), Simmons and Ropp7 measured densities of (only) saturated LiClO4 solutions at five temperatures from 273 to 303 K, and Geffcken8 at six concentrations over the range 2.0 ≤ m/mol·kg−1 ≤ 4.9 at 298 K. Subsequently, Jones9 represented his measurements of ρ(LiClO4(aq)) at concentrations up to ∼0.1 mol·L−1 at 298 K with an equation, and Haase and Dücker10 listed densities at 0.001 ≤ c/mol·L−1 ≤ 2.8 and nine temperatures from 273 K to 333 K. All of these studies were made by pycnometry (Haase and Dücker did not specify the method used but the large scatter in their data is consistent with this technique) and were of limited precision, with densities cited to 3,7 4,9,10 or (at best) 56,8 decimal places, with little or no information provided about the purity of the salt employed. © XXXX American Chemical Society
Although the data summarized above are of somewhat dubious quality (even at 298 K) they have been used by Söhnel and Novotný11 and by Aseyev and Zaytsev,12 in their respective compilations of physicochemical data of electrolyte solutions, to derive smoothed values of ρ(LiClO4(aq)) under conditions (0 ≤ m/mol·kg−1 ≤ msatd and 273 ≤ T/K ≤ 373) that go far beyond the published experimental matrix. The present experimental situation is summarized in Figure 1, which plots the apparent molar volumes (Vϕ) of LiClO4(aq) at 298.15 and 333.15 K, calculated from the available density data.5−10 Clearly, there is considerable uncertainty in the existing volumetric properties, especially at low and intermediate concentrations, even at 298.15 K. It follows that the smoothed values derived from these data by Söhnel and Novotný11 (which were adopted by Aseyev and Zaytsev, but assigned an extra significant figure!)12 must be viewed with caution. This is especially true at higher temperatures because the values proposed11,12 are based largely on the scattered data of Haase and Dücker10 (Figure 1) and, as noted above, extend well beyond their experimental range, especially with respect to temperature. A detailed numerical comparison of the present results with those listed by Aseyev and Zaytsev12 is given in the Results and Discussion section below. Given these uncertainties it is apparent that a reinvestigation of the volumetric properties of well-characterized LiClO4(aq) solutions using modern methodology is appropriate. Accordingly, the present paper reports measurements of LiClO4(aq) Received: June 26, 2015 Accepted: March 9, 2016
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DOI: 10.1021/acs.jced.5b00535 J. Chem. Eng. Data XXXX, XXX, XXX−XXX
Journal of Chemical & Engineering Data
Article
mass %, (theoretical: 33.68 mass % for LiClO4·3H2O) probably indicating incomplete drying. The perchlorate content was also determined in triplicate by precipitation of KClO4(s) from ethanol solution with potassium acetate.13 After being dried to constant mass at 125 °C the perchlorate content was found to be 61.70 ± 0.01 mass % (theoretical: 61.98 mass %), which is in quantitative agreement (within the uncertainty limits) with the presence of 0.30 mass % of adventitious water in the solid LiClO4·3H2O. Two stock solutions (4 and 5 mol·kg−1) were prepared by weight from high purity water (Ibis Technology system, Australia) and the synthesized LiClO4·3H2O, making appropriate allowance for the additional water in the solid. More dilute solutions were prepared by weight dilution of the stock solutions; buoyancy corrections were applied throughout. Operation of the Vibrating-Tube Densimeter. Vibrating-tube densimeters (vtd’s) are the most widely employed method nowadays for the measurement of solution densities as they provide a useful trade-off between accuracy and ease of measurement.14,15 Because the present type of apparatus has not been widely used for high-precision studies it is appropriate to briefly describe some of its characteristics and the measurement protocol adopted in consequence. Being an electromechanical device, the major limitation on the accuracy of any vtd is the slow drift with time of its vibration period (the measured quantity).15 This problem is largely overcome in the most accurate isothermal studies by bracketing solution measurements with immediate before-andafter measurements of the pure solvent, with appropriate interpolation. The apparatus used in the present study, an Anton Paar (Austria) model 5000 M vtd, has an inbuilt thermostat with a stated accuracy (at temperatures close to 298 K) of better than ±2 mK and a vibration-frequency stability corresponding to a few ppm (μg·cm−3) in density determinations. Of course the overall accuracy of the measured solution densities depends on many factors including calibration, solution preparation, and especially solute purity.14 The densimeter thermostat can be programmed so that densities are determined automatically over a specified temperature range. While this is convenient for the experimenter, it precludes the immediate before-and-after solvent measurements referred to above. Accordingly, the following experimental protocol was adopted. First the vtd was calibrated at (293.15, 313.15, and 333.15) K with laboratory air and high purity degassed water, using the air and water densities provided within the instrument software by the manufacturer. Note that these water densities differ slightly (by ∼4 μg·cm−3) from currently preferred values (IAPWS-95, see below). However, the former were used only to determine the densimeter constant (KB); because of the large difference in density between air and water such differences have a negligible effect (
