CK-12-Chemistry Intermediate

http://www.ck12.org Chapter 21. Acids and Bases

Substituting into the Kaexpression gives the following:

Ka= 5. 6 × 10 −^10 =

x^2

2. 00 −x

≈

x^2

2. 00

x= [H+] =

√

5. 6 × 10 −^10 ( 2. 00 ) = 3. 3 × 10 −^5 M

pH = -log(3.3× 10 −^5 ) = 4.48

A salt produced from a strong acid and a weak base yields a solution that is acidic.

Salts That Form Neutral Solutions

A salt that is derived from the reaction of a strong acid with a strong base forms a solution that has a pH of 7. An
example is sodium chloride, formed from the neutralization of HCl by NaOH. A solution of NaCl in water has no
acidic or basic properties, since neither ion is capable of reacting with water to produce any appreciable amounts
of H 3 O+or OH−. Other salts that form neutral solutions include potassium nitrate (KNO 3 ) and lithium bromide
(LiBr). The table below (Table21.12) summarizes how to determine the acidity or basicity of a salt solution.

TABLE21.12:pH of salts

Salt formed from: Salt Solution

Strong acid + Strong base Neutral

Strong acid + Weak base Acidic

Weak acid + Strong base Basic

Salts, formed from the reaction of a weak acid and a weak base, are more difficult to analyze because of competing
hydrolysis reactions between the cation and the anion. These salts will not be considered in this text.

Buffers

If only 1.0 mL of 0.10 M hydrochloric acid is added to 1.0 L of pure water, the pH drops drastically from 7.0 to 4.0.
This is a 1000-fold increase in the acidity of the solution. For many purposes, it is desirable to have a solution that
is capable of resisting large changes in pH when relatively small amounts of acid or base are added to them. Such a
solution is called a buffer. Abufferis a solution that contains significant amounts of both an acid and its conjugate
base. Both components must be present for the system to resist changes in pH. Commercial buffer solutions are
available with a wide variety of pH values.

One example of a buffer is a solution made of acetic acid (a weak acid) and sodium acetate (a salt containing its
conjugate base). The pH of a buffer consisting of 0.50 M CH 3 COOH and 0.50 M CH 3 COONa is 4.74. If 10.0 mL
of 1.0 M HCl is added to 1.0 L of the buffer, the pH only decreases to 4.73. This ability to “soak up” the additional
hydrogen ions from the HCl that was added is due to the reaction below.

CH 3 COO−(aq)+H+(aq)→CH 3 COOH(aq)

Since both the acetate ion and the acetic acid were already present in the buffer, the only thing that changes is the
ratio of one to the other. Small changes in that ratio have only very minor effects on the pH.

If 10.0 mL of 1.0 M NaOH were added to another 1.0 L of the same buffer, the pH would only increase to 4.76. In
this case the buffer takes up the additional hydroxide ions.

CH 3 COOH(aq)+OH−(aq)→CH 3 COO−(aq)+H 2 O(l)