General characteristics of nonelectrolytes solutions. Water solutions of nonelectrolytes. Acid-base characteristics of solutions, страница 4


Calculations  of values P, ∆tf,  ∆tb given above with the use of a molar concentration (C) only, molar concentration (Cm) or molal part (N2)  without taking electrolytic dissociation of substances into account and interactions of particles in a solution are true only for solutions of nonelectrolytes, that is compounds, not breaking apart on ions under influence of dissolvent (water).  The difference in calculated and experimentally found values P, ∆tf,  ∆tb for electrolytes due to increase of the total concentration of particles in a solution as a result of dissociation of molecules on ions.  Taking this difference into account van’t Hoff entered correction empiric factor, called now isotonic or van’t Hoff factor (i).  Isotonic factor is found  as the quotient between observable for electrolytes (experimental) and computational for nonelectrolytes (calculated) values of following sizes:

i = Pexp= ∆tfexp = ∆tbexp

          Pcal     ∆tfcal      ∆tbcal

For example, in a solution which contains 1 mol NaCI, freezing point lowers on 3,36° instead of 1,86° , that is almost twice as many as for analogic solutions of nonelectrolytes.

Later it was shown by Arrhenius works that the size of isotopic factor (i) depends on ion concentration.  According to the Arrhenius theory of electrolytic dissociation acids, base and salts are being exposed to dissociation of various degrees in solutions.  Arrhenius didn’t divide substances on weak and strong electrolytes.


The process of dissociation of weak electrolyte can be characterized by the degree of electrolytic dissociation and the constant of electrolytic dissociation. 



where ά – the degree of electrolytic dissociation;

           C1 – number of molecules, broken apart on ions;

           C – total number of dissolved molecules.

From this ratio follows that the degree of electrolytic dissociation ά can change from 0 to 1( or from 0% to 100%).  However, 100% of dissociation were never observed for weak electrolytes.  The degree of dissociation depends on the concentrations of the solution, temperature, nature of solute and dissolvent.

The process of electrolytic dissociation of weak electrolytes is  reversible.  Then, in according to the Le Châtelier’s principle, the dilution of the solution will remove the balance in the direction of increase of  electrolytic dissociation’s degree, that is a dissociation of weak electrolytes goes in diluted solutions better than in concentrated one.  With rise in temperature dissociation increases for electrolytes at which this process is accompanied by absorption of heat and decreases in case of allocation of heat.

The use of the law of active mass to electrolytic dissociation of weak electrolytes lets to get the expression for the constant of electrolytic dissociation. (K).

For example, in case of dissociation of acetic acid:

CH3COOH  rlarrow.gif (68 bytes) CH3COO- + H+

K=( H+) * (CH3COO-)   =1,8*10-5


From this balance we can see that increase in concentration of one of ions, for example, ( H+) or (CH3COO-)   will lead to a balance remove to the direction of formation of nondissociated molecules CH3COOH, that is the degree of dissociation of a weak electrolyte decreases greatly at entering a strong electrolyte with the ion of the same name in its solution.  So if in 0,1M solution CH3COOH add 0,1 mol CH3COONa then a dissociation of molecules of acetic acid will decrease in 360 times.

For weak electrolyte (like CH3COOH) having concentration C, the concentration of each of ions will be equal to άC, and a concentration of nondissociated molecules CH3COOH – (1 – ά);C.  Inserting these values in the equation of a constant, we shall receive:

K = άC * άC  =   ά2C2     = ά2C

        (1-ά)*C    (1-ά)*C      1-ά

For weak electrolyte the value (1-ά) can be accepted equal to 1 then K=ά2C

Then ά = √K/C

So, the degree of electrolytic dissociation inversely to a square root from the value of molar concentration of electrolyte.  In other words, when a weak electrolyte is diluted the degree of dissociation is increasing (the law of dilution by Ostwald).  The law of dilution is used for calculation of electrolytic dissociation’s constant of weak electrolytes defining the degree of electrolytic dissociation of electrolyte by experimental way in the solution of known concentration.  For example, will calculate KHBrO, if ά for 0,1M of solution HBrO equals to 10-4:

K = ά2C = (10-4) 2*10-1=10-9

At endless dilution of the solution the degree of electrolytic dissociation accepts limiting value and becomes equal to 1.  Will define the connection of isotonic factor with the degree of dissociation.  If before the beginning of dissociation there were C molecule in the solution  and from 1 molecule come υ ions at dissociation then the total number of particles (molecules and ions) after dissociation will be equal to

C* (1- ά) + υάC 

The quotient of the total number of particles in the solution of a weak electrolyte as a result of its dissociation to the number of particles

i= C(1- ά)+ υάC  =(1- ά) + υά=1+ ά(υ-1)


Where ά = i-1/υ-1, that is greater ά, greater i.