1. The first six acids in the table from the data booklet are strong acids because they all react quantitatively with water to form hydronium ions.

2. HF(aq) + H2O(l)  H3O+(aq) + F-(aq)
Most other acids are weak, in that they react with water to a much lesser extent.
For example, hydrofluoric acid:
HF(aq) + H2O(l)  H3O+(aq) + F-(aq)
undissociated molecules
dissociated molecules
The acid ionization constant (Ka) indicates the extent to which an acid will react with water. It is a ratio of the dissociated form of the acid to the undissociated form. The advantage of using Ka over % ionization to express strength is that once determined, it is valid over a wide range of concentrations, unlike % ionization which is concentration specific.

3. Calculating Ka from Amount Concentrations
There are two common calculations involving the Ka constant:
Calculating a Ka value from measured (empirical) concentration data.
Using a Ka value to predict the concentration of hydronium ions or the pH when the initial concentration of weak acid is known.
Calculating Ka from Amount Concentrations

6. Calculating [H3O+(aq)] from Ka
The Rule Of 1000
The value of x in the denominator can be omitted whenever the original concentration of the acid is at least 1000 times the numerical value of the Ka.
For any weak acid:
***Setting the concentration of the hydrogen ion to the concentration of the conjugate base is restricted to those cases where the initial concentration of the acid is MUCH larger (about 1000 times greater) than the Ka value. If this is not the case, then you must use the quadratic formula given in the data booklet to determine the concentration of the hydrogen ion.
No questions will be given to you that require you to use the quadratic formula. However, you may be asked if a given problem requires it. As a result, you should develop a habit of seeing if the assumption holds prior to using it!

9. Homework:
Read pgs. 737 – 742
pg. 743 Practice #’s 1 – 9

10. Base Strength and the Ionization Constant, Kb
All ionic hydroxides completely dissociate upon dissolving, so they are considered to be strong bases.

11. The Base Ionization Constant- Kb
Just like the with the acid ionization constant(Ka),the Kb value of a weak base is determined using an equilibrium law expression for a Bronsted-Lowry reaction between the weak base and water first and a substitution of concentration values second.
Example: What is the Kb formula for hydrazine N2H4(aq) ?
Unlike with Ka, there is not a list of given values on p.8-9 of the data booklet for Kb. However, by manipulating the mathematical relationship between Ka and Kb, Kb can be calculated. The base ionization constant (Kb) indicates the extent to which a base will react with water.
K w = Ka x Kb
Where Kw = 1.0 x
H2O(l) + N2H4(aq) OH-(aq) + N2H5 +(aq)
Kb = [OH-(aq)][ N2H5 +(aq)]
[N2H4(aq)]

12. The Ka–Kb Relationship for Conjugate Acid–Base Pairs
Consider the case of a general weak acid of the form HX(aq) and its reaction with water:
HX(aq) + H2O(l)  H3O+(aq) + X-(aq)

13. Ka × Kb = Kw X-(aq) + H2O(l)  HX(aq) + OH-(aq)
For the conjugate base X–(aq):
X-(aq) + H2O(l)  HX(aq) + OH-(aq)
Notice what happens when we multiply these two equilibrium expressions:
Ka × Kb = Kw
(recall: Kw = 1.0 × 10–14)

14. Calculating Kb from Amount Concentrations

15. Calculating [OH–(aq)] from Kb

17. The Rule Of 1000
The value of x in the denominator can be omitted whenever the original concentration of the base is at least 1000 times the numerical value of the Kb.
For any weak base:

19. The Effect of Amphoteric Entities
Remember, amphoteric entities can act as an acid or as a base.
To decide which one, compare the values of the Ka and Kb:
If Ka > Kb, then it acts as an acid.
If Ka < Kb, then it acts as a base.

21. Homework: Read pgs. 744 – 750 pg. 746 Practice #’s 10 – 13
pg. 750 Section 16.3 Questions #’s 1 – 10