Sodium enthalpy curves

Figure 1. Sodium dodecyl sulfate enthalpy curves at various temperatures...

Figure 2. Sodium dodecyl sulfonate and sodium decyl sulfonate enthalpy curves...

Figure 4. Sodium dodectjl sulfonate enthalpy curves in 5% n-butanol...

Figure 5. Sodium dodecyl sulfonate enthalpy curves at 35°C in alcohol backgrounds...

FIG. 15 Free enthalpy of adsorptioin on Na-illite and HDP-illite derivatives in methanol(l)-benzene(2) mixtures. Na sodium-illite curves 1-3 HDP-illites. 

Figure 5 Enthalpies of transfer of 2-butoxyethanol from water to sodium decanoate solutions at 25°C. Simulations (curves A and B) with a chemical equilibrium model.

The curved boundary lines on which the isotherms of Fig. 16.8 terminate represent conditions of temperature and concentration under which solid phases form. These are various solid hydrates of sodium hydroxide. The enthalpies of all single-phase solutions lie above this boundary line. The enthalpy-concentration diagram can also be extended to include solid phases. 

Consequently, the temperature dependence of /c( ) is given by the temperature dependence of the two rate constants k yc and k )s and of the equilibrium constant Kcs- The curved lines in Figure 18-3 thus show the transition from the more reactive, and at lower temperatures, more stable solvated ion pairs to the less reactive contact ion pairs. An Arrhenius equation with the preexponential factor A and activation energy E can be written for each rate constant, and a van t Hoff equation with the enthalpy AH and the entropy AS can be written for each equilibrium constant. Typical values for the polymerization of poly(styryl) sodium are given in Table 18-3. 

To relate now the solution enthalpy to the slope of solubility curves, in Table 3.3 the AsHoo values are compiled for selected salts included in Figure 3.17. The steep increase of the solubility curve of potassium nitrate correlates with its comparatively high-positive solution enthalpy. Consistently, the weak (positive) temperature dependence of the sodium chloride solubility is expressed by a low (also positive) AsHoo value. Negative solution enthalpies occur for anhydrates of salts forming stable hydrates at room temperature (such as sodium sulfate and sodium carbonate) where hydration is connected vhth a strongly exothermal effect. 

The relations between A//(sol) and the molar integral and differential enthalpies of solution are illustrated in Fig. 11.9 on the next page with data for the solution of crystalline sodium acetate in water. The curve shows A//(sol) as a function of fsou with fsoi defined as the amount of solute dissolved in one kilogram of water. Thus at any point along the curve, the molality is /Mb = soi/(l kg) and the ratio A//(sol)/ soi is the molar integral enthalpy of solution A//m(sol, /mb) for the solution process that produces solution of this molality. The slope of the curve is the molar differential enthalpy of solution ... 

It is of considerable interest to consider in greater detail plots of the enthalpy interaction parameters for the alkali chloride-magnesium chloride systems. These are shown in Fig. 8, taken from the work of Kleppa and McCarty. Note in particiilar the sharp dip in X near N qi2 0.33 for the systems containing KCl, RbCl, and CsCl. This correlates with the "anomalous partial entropy curves for these mixtures, and undoubtedly reflects the tendency of these systems to form the complex anion MgCl. Note also that the dip is completely absent in the silver chloride and lithium chloride system, and occurs only as a very broad minimum in sodium chloride-magnesium chloride. This is consistent with the nearly ideal entropy found in these systems. 

Figure 38 shows the results of demicellization experiments with the negatively charged surfactant sodium lauroyl-alaninate (SLA) at different temperatures (A. Blume, M, Ambiihl, and H. Watzke, unpublished results). It can be seen that the heat effect due to demicellization changes sign close to a temperature of 35°C. Similar curves are obtained for a variety of other surfactants. However, the temperature where A// = 0 varies with the headgroup stnicture of the surfactant. For anionic and cationic surfactants the temperature is usually between 20 and 30°C, whereas for non-ionic surfactants, such as octylglucoside (OG), it is close to 50 °C [123] The cmc can be easily and precisely detennined from the first derivative of the curves shown in Figure 38. The cmc is slightly temperature dependent and has a minimum where the demicellization enthalpy is zero.

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