Acids and Bases: An Introduction
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What are the H3O+ concentrations in the following solutions of strong acids and bases?

Acids

4.0 M HCl : 

 2.5 x 10-15 M 1 x 10-4 M 4.0 M 

0.01 M HNO3 : 

 0.01 M 1 x 10-12 M  1 M

Bases

6.0 M NaOH : 

 1.7 x 10-15 M 1 x 10-8 M 6.0 M 

1 x 10-5 M KOH :

 1 x 10-5 M 1 x 10-9 M 9.0 M 

Remember that a strong acid completely ionizes so a 1 M solution of a strong acid will have an H3O+ concentration of 1 M.

Remember that a strong base completely ionizes so a 5 M solution of strong base will have an OH- concentration of 5 M. In any aqueous solution:

[H3O+] [OH-] = Kw = 1 x 10-14

Good! As you can see from the above examples, the H3O+ concentration of aqueous solutions can have a huge range of values.







In 1909 a biochemist suggested a simpler way of expressing the acidity of a solution: pH. The pH of a solution is the negative logarithm of the H3O+ concentration:

pH = - log [H3O+]

What are the pHs of the above solutions?

[H3O+]

pH

0.1 M HCl

0.1 M

0.1

1

-1

0.001 M HNO3

0.001 M 

0.001 

1 

3

0.01 M NaOH 

1 x 10-12

.01 

2 

12 

1 x 10-5 M KOH

1 x 10-9

1 x 10-5

5 

9 

Remember that pH is equal to the negative log of the H3O+ concentration.

pH = - log [H3O+]

Good! While it may be difficult now to see how using pH simplifies things, you can see that the acidity of these solutions lie in a range of 11 pH units whereas before, the H3O+ concentrations were between 1.7 x 10-15 and 4 - a huge range!

Just as a solution has a pH, a solution also has a pOH. The pOH is equal to the negative logarithm of the OH- concentration:

pOH = - log[OH-]

Later on, you will see that this notation is used often when very small concentrations are measured. You may report the magnesium ion concentration as pMg2+ or the iron (III) concentration as pFe3+.