The Entropies of Ions in Aqueous Solution: I. Dependence on Charge and Radius

Laidler, Keith J.

doi:10.1139/v56-144

Cited by 36 publications

(12 citation statements)

References 8 publications

(9 reference statements)

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Order By: Relevance

“…This phenomenon had yet been observed when the complex strength stems from the gain of order [8]. For a given ligand, the Δ r 0 An(C 2 O 4 ) i n−2i (aq) values tend to increase as the effective charge of the actinide ion increases, as expected for electrostatic bonding [43]. In this study, for each complexation (except the first one), the expected following entropic reaction order is observed: Δ r .…”

Section: Discussionsupporting

confidence: 53%

Actinide oxalate complexes formation as a function of temperature by capillary electrophoresis coupled with inductively coupled plasma mass spectrometry

et al. 2014

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confirmed the formation of the 1 : 1, 1 : 2 and 1 : 3 complexes for U(VI) and Am(III); the formation of the 1 : 1 and 1 : 2 complexes for Np(V) and the formation of only 1 complex for Pu(V). For each complex, the thermodynamic parameters (the Gibbs energy Δ r ( ), the molar entropy change Δ r ( ) and the molar enthalpy change Δ r ( 0 )) were fitted to the experimental data. The effect of the ionic medium was treated using the specific ion interaction theory and the thermodynamic parameters at zero ionic strength were compared to previously published data.

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Section: Discussionsupporting

confidence: 53%

Actinide oxalate complexes formation as a function of temperature by capillary electrophoresis coupled with inductively coupled plasma mass spectrometry

et al. 2014

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“…This has been described as leading to the formation of icebergs about the solute but it is probably a better approximation to specify t h a t a particular water ~nolecule in the primary solvation shell spends Inore time in each configuration than is characteristic of bull; water (24,25). T h e net effect of this loss of freed0111 will clearly be proportional to the n~l m b e r of water nlolecules in the surface layer thus leading to such regularities between structures and solvation paranleters a s are recognized in the Barklay-Butler rule (26).…”

Section: Aqueous Solutions Of Non-electrolytes and Ionsmentioning

confidence: 99%

Solvolysis in Hydrogen and Deuterium Oxide: Iii. Alkyl Halides

Laughton¹,

Robertson²

1959

Can. J. Chem.

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Rate data for the hydrolysis of a series of a-ancl 8-methylated ally], benzyl, and cycloalliyl halides in light and heavy water are compared. The major factor determining the rate ratio ko o/kn,o appears to be the relative stability of the initial solvation shell. \Vhen the relative vi2cosity of Hz0 and DzO is used as a measure of the relative stability of structure in bulk solvent, the observed kinetic isotope effect and the temperature dependence of the isotope effect can be ratio~ialized in terms of accepted properties of aqueous solutions.Attention is drawn to the a~iomalous rate ratio knn,/knr > 1 which appears to be characteristic of solvolysis of these alliyl halides in pure water in contrast to other media.In an extension of recent solvolytic studies of strongly solvated substrates, chiefly sulphonates ( I ) , rates were obtained for the solvolysis of a variety of halogen compounds in light and heavy water. This work was done on the premise t h a t the isotope effect derived froin the rate ratio kn?o/krI?o was insensitive to temperature, a conclusion based on earlier work on the benzenesulphonates and inetl~anesulpl~onates (2). However, a recent determination of AC? for isopropyl bromide in light and heavy water (3) has cast doubt on the general applicability of this assumption. Since it was not practical to determine the AG: for each of the alltyl halides for which we had a rate ratio a t one temperature, we give here the rate ratios a t the convenient experimental temperatures, both because we propose t o refer to them later and because even subject to the above proviso they provide, for the detailed ~nechanism of hyclrolysis, clues which deserve cliscussion.T h e methods of ineasuring ltinetic rates were siinilar to those used previously (4). Easily volatile solutes were added under vacuum to the reaction flask, containing deaerated water and the potassium salt of the corresponding acid (0.003 M) either from a gas burette or from an evacuated side arm. With some solutes of low solubility, especially the benzyl bromicles, a saturated solution was formed, and a fritted glass dislc in the filling line was required to remove the sinall undissolved globules of the halide. With these latter solutes, supersaturation difficulties were occasionally encountered, a s shown by non-linear rate plots.YIost of the halides were commercial samples (chiefly Eastnlan I

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“…Nevertheless the continuum model has met with much criticism since, for example, according to Powell and Latimer (1951), Cobble (1953), and Connick and Powell (1953), it cannot be used to calculate the hydration energies of other than monovalent ions. Laidler (1956) does not agree with their argument. In any case it is clear that this mathematical model is fairly crude and that a refinement should take some account of the discrete nature of the water molecules.…”

Section: Introduotionmentioning

confidence: 85%

The Interaction Energy of an Isolated Monovalent Ion

Ross¹

1968

Aust. J. Phys.

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The interaction energy of a monovalent ion in an a.queous medium at 25°C is determined. It is also found that the water molecules in the first hydration shell of the ion have a mean dipole moment far in excess of their permanent dipole moments. Thus, for example, the increase in the dipole moment of the attached water molecules. due to the presence of an ion is about 60% for the small four-coordinated Li+ ion and about 30% for the larger four-coordinated 1-ion. Calculations are also carried out on the assumption that the ions are six coordinated. I. INTRODUOTIONA formula for the Gibb's free energy of hydration of an isolated ion in an aqueous solution was first obtained by Born (1920). He assumed that the ion could be replaced by an empty spherical cavity containing a point charge at its centre with the surrounding medium having the macroscopic dielectric constant of water. This model has been used by many authors, including Webb (1926), Latimer, Pitzer, and Slansky (1939), and Noyes (1962, to mention but a few. Nevertheless the continuum model has met with much criticism since, for example, according to Powell and Latimer (1951), Cobble (1953), andConnick andPowell (1953), it cannot be used to calculate the hydration energies of other than monovalent ions. Laidler (1956) does not agree with their argument. In any case it is clear that this mathematical model is fairly crude and that a refinement should take some account of the discrete nature of the water molecules. It is with this fact in mind that we shall develop a discrete model here.As a guide we can refer to the papers of Bernal and Fowler (1933), Verwey (1941, 1942, Rowlinson (1951), Buckingham (1957), Vaslow (1963), Ross (1968a, and Ross and Levine (1968). All but the last two papers omitted the polarizability ofthe water molecules, and in each case the water molecule was replaced by a point dipole (and sometimes 1\ point quadrupole as well). In fact Ross (1968a) showed that, at 25°C, a good approximation can be obtained by replacing the polarizable water dipoles of the first hydration shell in the position of minimum electrostatic energy with the water outside this shell replaced by a dielectric continuum. Now Bernal and Fowler (1933) indicated that the most likely coordination number oia monovalent ion is four, whereas Verwey (1941Verwey ( , 1942, who took into consideration the geometrical constraints in the second hydration shell, suggested that it is six or even eight. In the present paper we shall deal only with coordination number four or six, the former being the most probable. Similar calculations were

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The Entropies of Ions in Aqueous Solution: I. Dependence on Charge and Radius

Cited by 36 publications

References 8 publications

Actinide oxalate complexes formation as a function of temperature by capillary electrophoresis coupled with inductively coupled plasma mass spectrometry

Actinide oxalate complexes formation as a function of temperature by capillary electrophoresis coupled with inductively coupled plasma mass spectrometry

Solvolysis in Hydrogen and Deuterium Oxide: Iii. Alkyl Halides

The Interaction Energy of an Isolated Monovalent Ion

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