Molecular structure and acid Strength/Common Ion effect

Identify factors that affect acid strength

The stronger an acid, the more it will dissociate.   From the Ka equations, this means it will be more product heavy and this means it will be the Ka values that are larger. 

To predict strength for oxy-acids, it helps to remember that the definition of a Bronstead acid is a proton donator.  An acid will readily donate a proton the more the anion it leaves behind is stable. 

There is an increase in strength with increasing oxygen atoms since O is very electronegative.

Another way to look at it: 

For sulfuric acid (H₂SO₄), if you draw the Lewis structures it is clear that there are more resonant structures.   Since there are more, this means that the anion charge that is formed when it donates its first proton is more distributed and thus more stable.  The Ka value for the first ionization of sulfuric acid is 

With one less oxygen, you have sulfurous acid, H₂SO₃. There are less oxygens to resonate the negative charge and thus, the negative charge remains on the oxygens longer which is more unstable (see below).  Sulfurous acid is a weak acid with a lower equilibrium constant.  

For Binary acids, it once again depends on how easily it loses the proton.  Although HI and HCl are both strong acids, HI is stronger.  The atomic radius of iodine is bigger, and the bond with hydrogen is weak.  Thus, it dissociates easily, making it a better proton donor or a stronger acid. 

Demonstrate the common-ion effect: a solution of a weak acid and a strong acid

If you have a weak acid in the presence of a small acid, you must take into consideration the amount of hydronium ion that is present from the strong acid in your final calculation:

Example:

Find [H⁺] in a solution 5.00M HNO₂ and 0.01 M HNO₃. The K_a for HNO₂ is 5.6 x 10^-4.

Solution: 

x = 0.28 = [H⁺]

pH = -log [H⁺]

pH = 0.55

Demonstrate the common-ion effect: a solution of a weak acid and a salt of that weak acid

A common ion effect of an acid can be introduced when the salt of a weak acid is in the same solution as a weak acid.   You must account for the present of the reactant and product in your initial concentrations.

Conveniently for weak acids, these calculations can be worked out more simply using the Henderson Hasselbach equation:

where pKa = – log (K_a)

[A^–] = the concentration of the salt 

[HA] = the initial concentration of the weak acid

Example: 

Find [H⁺] in a solution 1.0M HNO₂ and 0.225M NaNO₂. The K_a for HNO₂ is 5.6 x 10^-4.

Solution:

This can be set up as an ICE table, but more easily you can set this up as a Henderson Hasselback equation: