Comments re Chapter 11 Problems:

11-2

Another T/F exercise. The points made are subtle, but very important. In part (b) of this problem it is important to know that the activity coefficient calculated from the Davies equation is the same for the cation and its associated anion, but that the molality of the cation and the molality of the anion must each be multiplied by that activity coefficient. Part (d) is rather subtle. Remember that if a very small amount of HCl is added to water so that the molality is 1.0 x 108 mol/kg then the solution is acidic and the pH < 7.00.

11-6

Exercise in calculating ionization constants from data in thermodynamic tables. Note that very small and very large equilibrium constants are seldom known very well.

11-8

A quick calculation to show that the log10mH+ of a NaCl(aq) solution is not 7.00, even if the temperature is 298 K. Neither the Na+(aq) ion nor the Cl-(aq) undergoes hydrolysis, but they change the ionic strength, which lowers the activity coefficients of the H+(aq) and OH-(aq) ions and so increases their concentrations.
NB/ It is true that -log10aH+ of aqueous solutions at 298 K is 7.00. The pH is defined as -log10aH+ rather than as -log10mH+ or as -log10[H+].

11-9

It is important to remember that Kw changes with temperature; the only temperature at which the value is 1.0 x 10-14 is 298 K.
(a) Use an empirical expression to get Kw at normal body temperature.
(b) Estimate Kw at normal body temperature by using the expression for the temperature dependence of ln K that was introduced earlier in the semester.

 

 

11-9

It is important to remember that Kw changes with temperature; the only temperature at which the value is 1.0 x 10-14 is 298 K.
(a) Use an empirical expression to get Kw at normal body temperature.
(b) Estimate Kw at normal body temperature by using the expression for the temperature dependence of ln K that was introduced earlier in the semester.

11-4

(Not a short problem, but very important).
All parts require solving an equilibrium problem for a reaction involving ions. Solutions containing ions are not ideal; activity coefficients for uncharged molecules are usually assumed to be 1, but activity coefficients for ions must be determined experimentally or estimated using an equation like the Davies equation.
(a) Dissociation of formic acid in pure water.
(b) Dissociation of formic acid in a solution containing an electrolyte that changes the ionic strength but that does not participate in the chemical reaction being considered.
(c) Dissociation of formic acid in a solution that also contains the formate ion, which (along with the accompanying Na+ ion) changes the ionic strength, and which also participates in the chemical reaction being considered (common ion effect). Note that this third solution is a buffer because the concentrations of the formic acid (undissociated acid) and formate ion (conjugate base) are similar.

See how including the activity coefficient for the uncharged, undissociated acid affects the calculations. See problem 11.36 for the equation.

(extra)

Calculate the pH = - log10aH+ for a 1.00 x 10-8 m HCl(aq) solution at 298 K.
Assume complete dissociation of the HCl. Use Kw = 1.00 x 10-14.
Use activity coefficients or show that they are not necessary.
NB/ Solution by the method of successive approximations may not work well so it may be better to solve by some more exact method.

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