Enthalpy of Solution and Hydration

4 Key excerpts on "Enthalpy of Solution and Hydration"

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  eBook - PDF

  Solvent Effects on Chemical Phenomena

  • Edward Amis(Author)
  • 2012(Publication Date)
  • Academic Press

    (Publisher)

  Relative values of the ionic entropies are recorded. These relative ionic entropies are referred to ^298.i of H + taken as zero. The most direct method of obtaining ionic entropies in solution is to sum the AS of solution and the entropy of the solid salt. The AS of solution is calculated from the free energy and enthalpy of solution using the equation AG = AH-TAS (3.75) where AG and AH are the free energy and enthalpy of solution, respectively. In recording the relative entropies referred to H + taken a zero, rather than these sums are frequently given. It has been pointed out [246, 248] that the so-called entropy of hydration is a function of the size and charge of the ion: *S hydmion =f(e 2 /r) (3.76) where e is the charge and r the radius of the ion. The entropy of hydration is defined as the difference between the partial molal entropy of the ion in solution and the entropy of the ion in the gaseous state. This latter entropy, provided it is all translation, can be calculated from the Sackur equation ^298.i = iR In M + 26.03 (3.77) where R is the gas constant in calories/mole/degree, M is the molecular weight, and 26.03 is the constant for the gas at 1 atm and 298.1°K. From straight-line plots of ionic entropies versus the reciprocals of the ionic Ill METHODS OF MEASUREMENT 93 radii as given by Pauling [250], it was concluded [248] that specific hydration effects are small compared to the electrostatic action of the charge on the water dipoles. In fact it was concluded earlier [246] that specific hydration effects do not exist, and that the chemical properties of these solutions of ions are those simply of a charged sphere of a given size in a medium of a certain dielectric constant. By ionic entropy measurements [251, 252] the hydration number of hydro-gen ion was found to be that given in Table 3.6. Values of the hydration number of various ions determined from ionic entropies have been listed [190, 253] and are given in Table 3.6.

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  Chemical Thermodynamics

  Theory and Applications

  • W.J. Rankin(Author)
  • 2019(Publication Date)
  • CRC Press

    (Publisher)

  12   Electrolyte solutions

  Scope

  This chapter examines the thermodynamics of solutions containing ions and shows how the concepts of enthalpy, entropy, Gibbs energy and activity are applied to ionic species in solutions.

  Learning objectives

  1. 1. Understand the nature of electrolytes and electrolyte solutions, in particular aqueous solutions.
  2. 2. Understand the methodology by which values can be measured and assigned to
     Δ f

     H 0
     ,
     Δ f

     S 0
     and
     Δ f

     G 0
     for ions in aqueous solutions.
  3. 3. Understand the concept of the activity of ions and of mean ionic activity.
  4. 4. Be able to estimate the activity of ions and electrolyte solutes in dilute solutions for fully dissociated and partially dissociated electrolytes.

  12.1 Introduction

  Chapter 9 examined the thermodynamics of solutions of neutral species. In this chapter we consider the thermodynamics of solutions containing charged species. The thermodynamic principles developed in Chapter 9 (chemical potential, activity, etc.) apply to these solutions, but their application is made complicated by the fact that the charged species interact strongly with each other. The charged species we are referring to are atoms or molecules that, through loss or gain of electrons, have a net electrical charge. These are called ions . Positively charged ions, with fewer electrons than protons, are called cations , while negatively charged ions, with more electrons than protons, are called anions . Compounds which dissociate into ions when they are dissolved in solutions, heated above a certain temperature or when they melt are called electrolytes * . The process by which a compound dissociates, partially or completely, into ions is called ionisation

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  Thermodynamics of Geothermal Fluids

  • Andri Stefánsson, Thomas Driesner, Pascale Bénézeth(Authors)
  • 2018(Publication Date)
  • De Gruyter

    (Publisher)

  This equation of state has been calibrated for 21 aqueous species (Holland and Powell 1998, 2011) but its accuracy and applications have not yet been tested. Recently, Dolejs and Manning (2010) evaluated mineral solubilities in aqueous fluids up to very high temperatures and pressures by a modified density model, which is internally con-sistent with hydration energetics and should provide a foundation for a new equation state for aqueous species. Their approach is based on two heuristic observations: (i) the intrinsic volu-metric properties of unhydrated species are closely approximated by those of the corresponding solid phase, and (ii) the caloric hydration properties become a simple expansion of enthalpy, entropy and heat capacity when evaluated at constant solvent density. This allows the intrinsic and hydration properties to be described by two distinct pressure-temperature paths, which leads to a particularly simple, low-parameter equation of state. The standard Gibbs energy of dissolution (A ds G) is conventionally defined as: A, s G = G aq -G s< (108) Expanding the standard Gibbs energies for both standard states in temperature and pressure yields: A

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  Chemistry, 5th Edition

  • Allan Blackman, Steven E. Bottle, Siegbert Schmid, Mauro Mocerino, Uta Wille(Authors)
  • 2022(Publication Date)
  • Wiley

    (Publisher)

  Notice, however, that when ‘theory’ predicts relatively large enthalpies of solution, the experimental values are also relatively large and both values have the same sign (except for NaCl, where the values are close to 0 and fall either side of it). Notice also that the changes in values show the same trends when we compare the three chloride salts — LiCl, NaCl and KCl — and the three bromide salts — LiBr, NaBr and KBr. Temperature can have a significant effect on the solubility of a solid solute in a liquid. Figure 10.16 shows a plot of solubility versus temperature for a number of ionic salts; it is obvious that all but one 466 Chemistry of these examples become more soluble as the temperature increases, which we might intuitively expect. There are few salts that become less soluble as the temperature increases. TABLE 10.2 Lattice enthalpies, hydration enthalpies and enthalpies of solution for some group 1 metal halides at 25 °C Δ sol H (a) Compound Lattice enthalpy (kJ mol -1 ) Hydration enthalpy (kJ mol -1 ) Calculated Δ sol H (kJ mol -1 ) (b) Measured Δ sol H (kJ mol -1 ) LiCl +833 -883 -50 -37.0 NaCl +766 -770 -4 +3.9 KCl +690 -686 +4 +17.2 LiBr +787 -854 -67 -49.0 NaBr +728 -741 -13 -0.602 KBr +665 -657 +8 +19.9 KI +632 -619 +13 +20.33 (a) Enthalpies of solution refer to the formation of extremely dilute solutions. (b) Calculated Δ sol H = lattice enthalpy + hydration enthalpy. FIGURE 10.15 Enthalpy of solution — the formation of aqueous sodium bromide pure water 1 mol NaBr(s) Enthalpy of system Δ sol H = –13 kJ mol –1 (The process is exothermic.) dilute solution of 1 mol NaBr Na + (aq) + Br – (aq) pure water Na + (g) + Br – (g) +728 kJ mol –1 (absorbed) to overcome lattice enthalpy (step 1) –741 kJ mol –1 (released) from dissolution (hydration enthalpy) (step 2) However, not everything is as straightforward as we would wish.

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