Freezing point depression electrolytes

The presence of a solute lowers the freezing point of a solvent if the solute is nonvolatile, the boiling point is also raised. The freezing-point depression can be used to calculate the molar mass of the solute. If the solute is an electrolyte, the extent of its dissociation, protonation, or deprotonation must also be taken into account. 

On the Theory of Electrolytes. I. Freezing Point Depression and Related Phenomena. (Zur Theorie der Elektrolyte. I. Gefrierpunktsemiedrigung und verwandte Erscheinun-gen). P. Debye and E. Huckel (Submitted February 27,1923), Translated from Physikalis-che Zeilschrift, Vol. 24, No. 9, 1923, pages 185-206, Classic Papers from the History of... 

ATt is the number of degrees that the freezing point has been lowered (the difference in the freezing point of the pure solvent and the solution). Kt is the freezing-point depression constant (a constant of the individual solvent). The molality (m) is the molality of the solute, and i is the van t Hoff factor, which is the ratio of the number of moles of particles released into solution per mole of solute dissolved. For a nonelectrolyte such as sucrose, the van t Hoff factor would be 1. For an electrolyte such as sodium sulfate, you must take into consideration that if 1 mol of Na2S04 dissolves, 3 mol of particles would result (2 mol Na+, 1 mol SO) ). Therefore, the van t Hoff factor should be 3. However, because sometimes there is a pairing of ions in solution the observed van t Hoff factor is slightly less. The more dilute the solution, the closer the observed van t Hoff factor should be to the expected one. 

A—Freezing-point depression is a colligative property, which depends on the number of particles present. The solution with the greatest concentration of particles will have the greatest depression. The concentration of particles in E (a non electrolyte) is 0.10 m. All other answers are strong electrolytes, and the concentration of particles in these may be calculated by multiplying the concentration by the van t Hoff factor. 

Salt is a strong electrolyte that produces two ions, Na+ and Cl, when it dissociates in water. Why is this important to consider when calculating the colligative property of freezing point depression ... 

Activity data for electrolytes usually are obtained by one or more of three independent experimental methods measurement of the potentials of electrochemical cells, measurement of the solubility, and measurement of the properties of the solvent, such as vapor pressure, freezing point depression, boiling point elevation, and osmotic pressure. All these solvent properties may be subsumed under the rubric colligative properties. 

It can be observed that g is the ratio between the observed osmotic pressure and the osmotic pressure that would be observed for a completely dissociated electrolyte that follows Henry s law [see Equation (15.47)], hence the name, osmotic coefficient. A similar result can be obtained for the boiling point elevation, the freezing point depression, and the vapor pressure lowering. 

A unitless correction factor that relates the relative activity of a substance to the quantity of the substance in a mixture. Activity coefficients are frequently determined by emf (electromotive force) or freezing-point depression measurements. At infinite dilution, the activity coefficient equals 1.00. Activity coefficients for electrolytes can vary significantly depending upon the concentration of the electrolyte. Activity coefficients can exceed values of 1.00. For example, a 4.0 molal HCl solution has a coefficient of 1.76 and a 4.0 molal Li Cl has a value of... 

Electrolytes, depending upon their strength, dissociate to a greater or less extenl in polar solvents. The extent to which a weak electrolyte dissociates may be determined by electrical conductance, electromotive force, and freezing point depression methods. The electrical conductance method is the most used because of its accuracy and simplicity. Arrhenius proposed that the degree of dissociation, a. of a weak electrolyte at any concentration in solution could be found from the rutio of the equivalent conductance. A. of the electrolyte at the concentration in question to (he equivalent conductance at infinite dilution A0 of the electrolyte. Thus... 

This procedure and the / V I Hall-Heroult process discussed below are examples of commercial uses of a colligative property (freezing point depression) to enable electrolytic reactions to be carried out more economically. 

Just as we discussed in Chapter 9, we can use measured activities of solvents (determined from vapor pressure, freezing-point depression, boiling-point elevation, or osmotic pressure) to determine activity coefficients of electrolytes in solution. For an ionic substance, the Gibbs-Duhem equation is... 

There are many measurement techniques for activity coefficients. These include measuring the colligative property (osmotic coefficients) relationship, the junction potentials, the freezing point depression, or deviations from ideal solution theory of only one electrolyte. The osmotic coefficient method presented here can be used to determine activity coefficients of a 1 1 electrolyte in water. A vapor pressure osmometer (i.e., dew point osmometer) measures vapor pressure depression. 

It has been suggested, for example, that the addition of ethylene glycol and polyethylene glycols enhances the F/T stability of polymer emulsions (2). Further, the addition of electrolytes has been proposed to achieve the same goal (5). All are freezing-point depressants and may influence the other factors as well. 

Partial Molar Volume 1 0. Cryoscopic Determination of Molar Mass 1 1. Freezing-Point Depression of Strong and Weak Electrolytes 12. Chemical Equilibrium in Solution... 

Freezing-Point Depression of Strong and Weak Electrolytes... 

In this experiment, the freezing-point depression of aqueous solutions is used to determine the degree of dissociation of a weak electrolyte and to study the deviation from ideal behavior that occurs with a strong electrolyte. 

Why is the observed freezing-point depression for electrolyte solutions sometimes less than the calculated value Is the error greater for concentrated or dilute solutions ... 

The Osmotic Coefficient.—Instead of calculating activity coefficients from freezing-point and other so-called osmotic measurements, the data may be used directly to test the validity of the Debye-Hiickel treatment. If 6 is the depression of the freezing point of a solution of molality m of an electrolyte which dissociates into v ions, and X is the molal freezing-point depression, viz., 1.858° for water, a quantity , called the osmotic coefficient, may be defined by the expression... 

A good deal of confusion and much controversy between early workers 54, 55, 77) resulted from a lack of understanding of the nature and extent of the self-dissociation of sulfuric acid, which is repressed by most solutes and thus affects the freezing-point depressions that they produce. Following Hammett and Deyrup 56) it became common practice to carry out cryoscopic measurements in sulfuric acid containing sufficient water to depress the freezing point to approximately 10.0° in order to repress largely the solvent self-dissociation. This is not, however, an entirely satisfactory procedure. The self-dissociation is not completely repressed (J), and allowance for it should still in principle be made. Moreover water has been found, in the case of nonelectrolytes and weak electrolytes at least, to have an effect on the depression... 

As we have emphasized, colligative properties depend on the number of solute particles in a given mass of solvent. A 0.100 molal aqueous solution of a covalent compound that does not ionize gives a freezing point depression of 0.186°C. If dissociation were complete, 0.100 m KBr would have an ejfective molality of 0.200 m (i.e., 0.100 m K+ + 0.100 m Br ). So we might predict that a 0.100 molal solution of this 1 1 strong electrolyte would have a freezing point depression of 2 X 0.186°C, or 0.372°C. In fact, the observed depression is only 0.349°C. This value for ATf is about 6% less than we would expect for an effective molarity of 0.200 m. 

Table 14-3 lists actual and ideal values of i for solutions of some strong electrolytes, based on measurements of freezing point depressions. 

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