Video 3: Buffers Equations

Video 2: Buffers

Video: Buffer introducation

Videos on Solutions:

for bonus, summarize the 5 key points.

Ch 19: Problems

(40) Key Concept What type of salt produces an acidic solution? A basic solution?Hint
(41) Key Concept What substances are combined to make a buffer? Hint
(42)Which of these salts would form an acidic aqueous solution?
a. KC2H3O2
b. LiCl
c. NaHCO3
d. (NH4)2SO4
(43)Using equations, show what happens when acid is added to an ammonium ion–ammonia buffer. What happens when base is added?

Ch 19: Intro to Buffers 2

Buffers
The addition of 10 mL of 0.10M sodium hydroxide to 1 L of pure water increases the pH by 4.0 pH units (from 7.0 to 11.0). A solution containing 0.20 mol/L each of ethanoic acid and sodium ethanoate has a pH of 4.76. When moderate amounts of either acid or base are added to this solution, however, the pH changes little. The addition of 10 mL of 0.10M sodium hydroxide to 1 L of this solution, for example, increases the pH by only 0.01 pH unit, from 4.76 to 4.77. Figure 19.27 shows what happens when 1.0 mL of 0.01M HCl solution is added to an unbuffered solution.


Figure 19.27 A buffer is a solution in which the pH remains relatively constant. a. The indicator shows that the buffered solution on the left and the unbuffered solution on the right are basic—pH about 8. b. After the addition of 1.0 mL of 0.01M HCl solution, the pH of the buffered solution shows no visible change. The pH of the unbuffered solution, however, is now about 3—the solution is acidic.
The solution of ethanoic acid and sodium ethanoate is an example of a typical buffer. A buffer is a solution in which the pH remains relatively constant when small amounts of acid or base are added. A buffer is a solution of a weak acid and one of its salts, or a solution of a weak base and one of its salts.
A buffer solution is better able to resist drastic changes in pH than is pure water. Figure 19.28 illustrates how a buffer works. Ethanoic acid (CH3COOH) and its anion (CH3COO−) act as reservoirs of neutralizing power. They react with any hydroxide ions or hydrogen ions added to the solution. For example, consider the buffer solution in which the sodium ethanoate (CH3COONa) is completely ionized.

Ch 19: Salt Hydrolysis

The ethanoate ion is a Brønsted-Lowry base, which means it is a hydrogen-ion acceptor. It establishes an equilibrium with water, forming electrically neutral ethanoic acid and negative hydroxide ions.

This process is called hydrolysis because it splits a hydrogen ion off a water molecule. The resulting solution contains a hydroxide-ion concentration greater than the hydrogen-ion concentration. Thus the solution is basic.
Ammonium chloride (NH4Cl) is the salt of a strong acid (hydrochloric acid, HCl) and a weak base (ammonia, NH3). It is completely ionized in solution.
NH4Cl(aq) → NH4+(aq) + Cl−(aq)

Figure 19.25 Vapors of the strong acid HCl(aq) and the weak base NH3(aq) combine to form the acidic white salt ammonium chloride (NH4Cl).
The ammonium ion (NH4+) is a strong enough acid to donate a hydrogen ion to a water molecule, although the equilibrium is strongly to the left.

This process is also called hydrolysis. It results in the formation of unionized ammonia and hydronium (hydrogen) ions. The [H3O+] is greater than the [OH−]. Thus, a solution of ammonium chloride is acidic. To determine if a salt solution is acidic or basic, remember the following rules:
Strong acid + Strong base → Neutral solution
Strong acid + Weak base → Acidic solution
Weak acid + Strong base → Basic solution

Ch 19: Intro to Buffers

Salt Hydrolysis
A salt consists of an anion from an acid and a cation from a base. It forms as a result of a neutralization reaction. Although solutions of many salts are neutral, some are acidic and others are basic. Solutions of sodium chloride and of potassium sulfate are neutral. A solution of ammonium chloride is acidic. A solution of sodium ethanoate (sodium acetate) is basic. Figure 19.24 shows a titration curve obtained by adding a solution of sodium hydroxide, a strong base, to a solution of ethanoic (acetic) acid, a weak acid. An aqueous solution of sodium ethanoate exists at the equivalence point.

The pH at the equivalence point is 8.7—basic.
For a strong acid–strong base titration, the pH at the equivalence point is 7, or neutral. This difference exists because some salts promote hydrolysis. In salt hydrolysis, the cations or anions of a dissociated salt remove hydrogen ions from or donate hydrogen ions to water. Depending on the direction of the hydrogen-ion transfer, solutions containing hydrolyzing salts may be either acidic or basic. Hydrolyzing salts are usually derived from a strong acid and a weak base, or from a weak acid and a strong base. Sodium carbonate, washing soda, is the salt of the strong base sodium hydroxide and carbonic acid, a weak acid. Ammonium nitrate, used in fertilizers, is the salt of the weak base ammonia and nitric acid, a strong acid. Soap is the salt of a strong base, usually sodium hydroxide, and stearic acid, a weak acid present in fats. In general, salts that produce acidic solutions contain positive ions that release protons to water. Salts that produce basic solutions contain negative ions that attract protons from water.
Sodium ethanoate (CH3COONa) is the salt of a weak acid (ethanoic acid, CH3COOH) and a strong base (sodium hydroxide, NaOH). In solution, the salt is completely ionized.

Ch 19: Acids/Bases 19.5 Buffers

Ch 19
Key Concepts
• When is the solution of a salt acidic or basic?
• What are the components of a buffer?
Vocabulary
• salt hydrolysis
• buffer
• buffer capacity

Bonus Question Ch 19 (EP 6)

Find Kb and pKb for the acetate ion CH3COO-. The ionization constant of CH3COOH is Ka=1.75 x 10-5 (this is to -5 power); Kw=1.00 x 10-14--this is to the -14 power.