Acids are bases are special types of chemicals. They are separated from other materials because of their properties.

As a general rule, an acid is any substance which releases or provides H⁺ ions. This H⁺ ion then reacts with H₂O to form the hydronium ion, H₃O⁺, because of the charge on the ion and the polar nature of water. The hydronium ion, H₃O⁺, is the reactive ion in solution. In order to save time, many chemists still refer only to the H⁺.

Complementary to acids are the compounds known as bases. Bases accept H⁺ ions from or release OH^- ions to the solution.

The Bronsted-Lowry model of acids and bases states the following:

    1) an acid is a proton donor
    2) a base is a proton acceptor
    3) in an acid-base reaction, a proton is transferred from the 
       acid to the base.

Consider that an acid releases a H⁺ ion into solution. Look at the H⁺ ion.

An H atom contains 1 proton in the nucleus and has a single electron orbiting that nucleus.

An H⁺ ion is an H atom that has lost an electron. An H atom that loses an electron is only a proton since there only was 1 proton and 1 electron at the start. Thus the H⁺ ion is a proton nd the terminology of Bronsted and Lowry uses this fact.

Acids and bases are classified according to the level to which they release H⁺ or OH^-.

Strong acids release all their H⁺ and are said to dissociate completely. The following example uses the generic term, "A", to be the anion of the acid.

          HA <-->  H+ + A-

          Ka = [H+][A-]

Strong bases release all their OH^- ions to solution and are also said to dissociate completely.

          MOH <-->  M+ + OH-

          Kb = [M][OH-]

As with the strong acid, a strong base has no MOH left so it is not included in the equation for the K.

Alternatively, the Bronsted-Lowry model shows the transfer of the H⁺ ion from one acid, HB, to another acid, HA.

          HB   +   A-    <-->    HA   +    B-
          |        |             |         |
  conjugate acid---|-------------|------conjugate base
               conjugate     conjugate
                 base          acid

When an acid and a base are both introduced to the same solution, the acid releases its H⁺ ion and the base releases its OH^- ion. The following reaction then occurs:

                 [H+][OH-]
             K = ---------
                   [H2O]

So we modify the K above by leaving out the [H₂O] since [H₂O] is constant. Our new constant is called K_w and is shown below.

     ACID                          NEUTRAL                        BASE
pH   0                               7                              14
     <-------------------------------|------------------------------->
[H+] 1                             10-7                           10-14
     <-----  acidity increases              basicity increases  ----->

     Strong acids         Weak acids    Weak bases        Strong bases

EXAMPLE 1:

A water sample is found to have a pH of 8.3. What is it's [H⁺]?

Multiply both sides by -1.

Take the antilog of both sides.

This is a mathematical identity.

So:

And in more convenient numbers:

EXAMPLE 2:

A water sample is found to have a pOH of 8.3. What is the pH of the sample?

Using the relationship: pH = 14 - pOH we find:

Multiply both sides by -1.

Take the antilog of both sides.

This is a mathematical identity.

So:

And in more convenient numbers:

Please note that a solution with a pH of 5.7 has nearly 400 times more [H⁺] than a solution with a pH of 8.3.

Recall that with strong acids the Equilibrium Constant, K_a, was:

Consider the dissociation of weak acid, HB, where B represents any anion.

Because the dissociation is partial, some HB remains. When we form the EQ constant, we get:

              [H+][B-]
         Ka = -------
               [HB]

The student should note from the text that any substance or ion that can cause the release of H⁺ ions to solution is considered an acid. If the K for the reaction is small (less than 0.01), the substance is considered a weak acid.

Example 3:

A solution of acetic acid is prepared. The acid concentration is 0.001 M. What is the pH of this solution?

The reaction:

The K for the reaction:

                             [H+][C2H3O2-]
                        Ka = ------------
                              [HC2H3O2]

Set up a table for initial and final conditions. Look up the K_a for acetic acid.
K_a = 1.8x10^-5

                      (x)(x)
               Ka = ----------- = 1.8 x 10-5
                      0.001-x 

                 (x)2
                -------  =  1.8 x 10-5
                 0.001

Then:

                x2 = 1.8 x 10-5 * 0.001

                x2 = 1.8 x 10-8

                  x = 1.34 x 10-4

                      (x)(x)
                 ---------------- = 1.8 x 10-5
                  0.001-0.000134 


And solve a second time:

                   (x)2
                 ------------ = 1.8 x 10-5
                   0.000866


                  x2 = 1.8 x 10-5 * 0.000866

                        x2 = 1.559 x 10-8

                        x = 1.25 x 10-4

The pH is:

When considering weak bases, the same pattern of logic holds except that the substance accepts H⁺ ions from the solution.

Here's a weab base ionization to consider:

             [NH4+][OH-]
        Kb = ----------
                [NH3]

To determine the pH one could easily find the pOH and subtract that from 14 to get pH.

So if you know K_a or K_b, you can calculate the other.

                Kw
        Ka = -------
                Kb

or

                Kw
        Kb = -------
                Ka