CHAPTER 15 - LECTURE 2:


15.5 Weak Acids And Acid Ionization Constants


Weak acids are only partially ionized in solution.
   There is a mixture of ions and non-ionized acid in solution.   That is, weak acids are in equilibrium:
               HA(aq) + H2O(l) ---> H3O+(aq) + A-(aq)
                                           <---
           or       HA(aq) --->  H+(aq) + A-(aq)
                                  <---
Since we use
H+ and H3O+ interchangeably:

   Ka = [H3O+][A-] / [HA]         or         Ka = [H+][A-] / [HA]
 

        Ka is the acid dissociation constant.

Note that [H2O] is omitted from the Ka expression;  H2O is a pure liquid, and its concentration is essentially constant.

The larger the K
a, the stronger the acid (i.e., the more ions
are present at equilibrium relative to the un-ionized molecules).

If Ka >> 1, then  the acid is completely ionized and the acid is a strong acid.   In general, a weak acid has a Ka of less than ~10-3.
 

Percent ionization:
Percent ionization is another method to assess acid strength.
All strong acids are 100% ionized!

7 strong acids: HCl, HBr, HI, HClO3, HClO4, HNO3, H2SO4
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Not so for weak acids.
       % ionization = ([H3O+]equil / [HA]0) x 100

 Percent ionization relates the equilibrium [H3O+] to the INITIAL [HA] concentration.
 The higher percent ionization, the stronger the acid.
 Percent ionization of a weak acid decreases as the molarity of the solution increases.
 For acetic acid, 0.05 M solution is 2.0 % ionized, whereas a 0.15 M solution is only 1.0 % ionized.
 

Calculate the % ionization from Ka:
 The Ka value for formic acid is 1.8 x 10-4.  Calculate the % ionization for a 0.120 M formic acid solution.    HCOOH is formic acid.

            HCOOH + H2O ----> H3O+ + HCOO-
                                    <----

    Ka= 1.8 x 10-4  =  ([H3O+][HCOO-])/[HCOOH]
                              = (x)(x)/(0.120-x)
            1.8 x 10-4  =  x2/0.120    (Assuming that x is small compared to 0.120)
            [1.8 x 10-4 (0.120)]1/2  = x
                      x = 0.00465 M = [H3O+]equil = [HCOO-]
    Therefore, % ionization = ([H3O+]equil / [HA]0) x 100
    % ionization = (0.00465M / 0.120M) x 100 = 3.88 %
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Calculating pH for Solutions of Weak Acids:
pH is the equilibrium concentration of H3O+.   By doing an equilibrium calculation using Ka , we can calculate [H3O+] at equilibrium: ([H3O+]equil).
Assume the initial [H3O+]0 and [A-]0 are zero; let x = change in moles of all species present, and  [HA]0 = initial conc. of acid

For a monoprotic weak acid, e.g., acetic acid, HC2H3O2:
                        Ka = [H3O+]equil[A-]equil / [HA]0 - x
                        Ka = (x)(x) / ([HA]0 - x) = x2 / ([HA]0 - x)
            solve for x ===>  [H3O+] ===> pH

HINT: If the value of x is small compared to the value of [HA]0 , we can assume that ([HA]0 - x) = [HA]0

This simplifies the quadratic and saves a lot of math!

Always check the final result when assuming x is small by plugging all values back into the equation!!!

FOR HELP SOLVING THESE KINDS OF PROBLEMS, check out             chapter 15Add2  !!!!

In general, if x < 5% of [HA]0, assume that     [HA]0 - x = [HA]0
You can often tell whether this will be true by comparing [HA]0 and Ka; if there is > 103 difference between [HA]0 and Ka, then assume that
([HA]
0 - x) = [HA]0

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15.6 Weak bases and Base Ionization Constants
 

Weak bases remove protons from substances.
There is an equilibrium between the base and the resulting ions; note that a base can be neutral, e.g., NH3,  or anionic, e.g., O2-

     (1)           B + H2O(l) ----> HB+(aq) + OH-(aq)
                                        <----

   or (2)        B- + H2O(l) ----> HB(aq) + OH-(aq)
                                         <----
 The base dissociation constant, Kb, is defined as
       (1)         Kb = [HB+][OH-]/[B]
     or (2)      Kb = [HB][OH-]/[B-]


Example:

          NH3(aq) + H2O(l) ---->  NH4+(aq) + OH-(aq)
                                        <----

 The base dissociation constant, Kb, is defined as
                      Kb = [NH4+][OH-] / [NH3]

 The larger the value of Kb, the stronger the base.   Bases generally have lone pairs or negative charges that attack protons, and weak bases often have a
 N atom
(substances called amines), e.g., pyridine C5H5N, methylamine NH2(CH3), dimethylamine NH(CH3)2, triethylamine N(C2H5)3

Anions of weak acids are also bases: e.g., OCl-

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