Chapter 10

Acids and Bases

Chapter 10 suggested problems: 38, 40, 42, 44, 46, 62, 64, 66, 68

Chapter Objectives:

After completing this chapter, you should, at a minimum, be able to do the following. This information can be found below and in your text. Sorry, no lecture notes for this material yet.

1. Correctly answer all of the questions suggested above and in the quiz for this chapter.
2. Define basic terms such as Arrhenius acid, Arrhenius base, Brønsted-Lowry acid, Brønsted-Lowry base, strong acid, weak acid, strong base, weak base, hydronium ion, conjugate acid, conjugate base, monporitc acid, diprotic acid, triprotic acid, polyprotic acid, amphoteric, amphiprotic, ion product constant for water, pH, pOH, neutralization, hydrolysis, buffer.
3. Describe which acids are Arrhenius acids and which are Brønsted-Lowry acids.
4. Describe which bases are Arrhenius acids and which are Brønsted-Lowry base.
5. Given a list of acids and bases, identify which are strong acids and bases and which are weak acids and bases.
6. Describe how water behaves as both an acid and a base. To what extent does this occur in a neutral solution of water at room temperature and pressure? Is it a product-favored or reactant-favored reaction?
7. Calculate the concentrations of strong acids and bases in aqueous solution.
8. Explain how the pH scale is based on water.
9. Calculate the pH of a solution of a strong acid or base.
10. Explain what neutralization and hydrolysis reactions are.
11. Explain how a buffer works.

Class Notes

1. General theory
1. Acid-base theories
1. Arrhenius
2. Brønsted-Lowry
3. Lewis - not discussed in text; ignore.

2. Strong and weak acids and bases (electrolytes)
1. Strong - dissociate abt. 100%
1. Acids: HCl, HBr, HI, HNO₃, H₂SO₄, HClO₄
2. Note: free ions do not exist in solution but are solvated by clusters of solvent molecules, so the presence of hydrogen ion in solution is approximated with H₃O⁺ (hydronium ion)
3. Bases: Group I and II hydroxides (Ca²⁺, Sr²⁺, Ba²⁺)

2. Weak - dissociate less than 100%, usually abt. 1-10%
1. Acids: organic acids (acetic acid HC₂H₃O₂), HF, H₂CO₃, H₃PO₄, H₂S;
2. All acids that are not on the strong acid list above are almost certainly weak acids
3. Bases: NH₃, ammonium hydroxide, amines;
4. All bases that are not on the strong base list above are almost certainly weak bases.

3. Note: the molecular formula for acids begin with "H"

3. Conjugate acid-base pairs
1. The generic behavior of any acid in aqueous solution is described by the following equation: HA + H₂O -> H₃O⁺ + A^-
2. The anion of the acid is called is conjugate base.
3. This is because the anion can react with water to form hydroxide ion: A^- + H₂O -> HA + OH^-
4. The generic behavior of any base in aqueous solution is described by the following equation: B + H₂O -> HB⁺ + OH^-
5. The cation of the base is called its conjugate acid.
6. This is because the cation can react with water to form hydronium ion: HB⁺ + H₂O -> H₃O⁺ + B
7. The equation in 4. does not do quite as good a job at describing base behavior as the acid equation in 1.; the conjugate acids of bases are not always cations
1. PO₄^3- + H₂O => HPO₄^2- + OH^-
2. HPO₄^2- + H₂O => PO₄^3- + H₃O⁺

4. Polyprotic acids and bases

2. The behavior of acids and bases in water
1. As stated previously, for acids: HA + H₂O -> H₃O⁺ + A^-
2. As stated previously, for bases: B + H₂O -> HB⁺ + OH^-
3. Examples
1. HNO_{3 (aq)} + H₂O_(l) -> H₃O⁺_(aq) + NO₃^-_(aq)
2. HC₂H₃O_{2 (aq)} + H₂O_(l) -> H₃O⁺_(aq) + C₂H₃O₂^-_(aq)
3. NaOH_(aq) + H₂O_(l) -> Na⁺_(aq) + OH^-_(aq) + H₂O_(l)
4. NH_{3 (aq)} + H₂O_(l) -> NH₄⁺_(aq) + OH^-_(aq)

4. The behavior of water: 2 H₂O_(l) -> H₃O^+_(aq) + OH^-_(aq)
1. 
2. 

5. Amphoteric, amphiprotic, and zwitterions
1. (source: http://www.psigate.ac.uk/newsite/reference/plambeck/chem1/p01155.htm)
   The term amphiprotic in modern acid-base chemistry is the replacement for the older term amphoteric. An amphiprotic substance is a substance which can act both as an acid and as a base because it contains at least one proton which can be given up and at least one site at which a proton can be acquired.
2. Water is amphiprotic
3. Most polyprotic acids have at least one amphiprotic ion. Using phosphoric acid as an example, the monohydrogen phosphate ion and the dihydrogen phosphate ion are both amphiprotic while phosphoric acid itself can only be an acid and the phosphate ion can only be a base.
4. Organic compounds which contain both a carboxylic acid group and an amine group on the same molecule are called amino acids. When an amino acid such as glycine, H₂NCH₂COOH, is dissolved in water. the carboxylic acid group loses a proton which is gained by the more basic amine group. This produces an ionic structure with opposite charges on both ends, a zwitterion. The zwitterion structure of glycine is +H₃NCH₂COO-. The protonated form of this amphiprotic zwitterion, +H₃NCH₂COOH, is the glycinium ion.
5. When an amphiprotic substance alone is dissolved in water, it will act both as an acid and as a base.

3. Concentration and pH
1. The concentration of strong acids and bases
1. What is the [H₃O⁺] of 0.25 M HNO₃? (0.25M)
2. What is the [H₃O⁺] of 0.25 M H₂SO₄? (0.50M)
3. What is the [H₃O⁺] of 0.25 M NaOH? (4 x 10^-14 M)
1. K_w = [H₃O⁺][OH^-] so [H₃O⁺] = K_w / [OH^-]

4. What is the [H₃O⁺] of 0.25 M Ba(OH)₂? (2 x 10^-14 M)

2. pH - used to denote acidity and basicity
1. The "p" function in chemistry
2. pH = - log [H₃O⁺]
3. pOH = - log [OH^-]
4. For water
1. [H₃O⁺] = [OH^-] = 1 x 10^-7 M
2. pH = 7, pOH = 7
3. pH + pOH = 14

5. The pH scale is based on water
1. Neutral: pH = 7, [H₃O⁺] = [OH^-]
2. Acidic: pH < 7, [H₃O⁺] > [OH^-]
3. Basic: pH > 7, [H₃O⁺] < [OH^-]

6. The pH scale is a logarithmic scale; pH = 5 is 10x more acidic than pH = 6 and 100x more acidic than pH = 7
7. Examples
1. What is the [H₃O⁺] of 0.25 M HNO₃? (0.602)
2. What is the [H₃O⁺] of 0.25 M H₂SO₄? (0.301)
3. What is the [H₃O⁺] of 0.25 M NaOH? (13.4)
4. What is the [H₃O⁺] of 0.25 M Ba(OH)₂? (13.7)

3. The pH of weak acids and bases
1. Since weak acids and bases do not dissociate 100%, calculating [H₃O⁺] depends on a knowledge of the equilibrium constant for the reaction HA + H₂O -> H₃O⁺ + A^-, or for bases, on a knowledge of the equilibrium constant for the reaction A^-+ H₂O -> HA + OH^-
2. K_a and K_b
3. What is the pH of a 0.15 M solution of acetic acid (K_a = 1.8 x 10^-5)?
4. Don't worry about calculating the pH of weak acids and bases unless you're given [H₃O⁺]

4. Neutralizations, salts, and hydrolysis
1. Neutralization: the reaction of an acid and a base to form water and a salt
2. Salts: ionic compounds formed during neutralization reactions in which the cation comes from a base and the anion comes from an acid
1. HCl_(aq) + NaOH_(aq) -> H₂O_(l) + NaCl_(aq)
2. H₂SO_{4 (aq)} + 2 NH₄OH_(aq) -> 2 H₂O_(l) + (NH₄)₂SO_{4 (aq)}
3. 2 HC₂H₃O_{2 (aq)} + Ca(OH)_{2 (aq)} -> 2 H₂O_(l) + Ca(C₂H₃O₂)_{2 (aq)}

3. Hydrolysis: the breaking of water into H⁺ and OH^-
1. The conjugate bases of acids cause hydrolysis
1. HA + H₂O -> H₃O⁺ + A^-
2. A^-+ H₂O -> HA + OH^-
3. The conjugate bases of acids are basic (i.e. form OH^-)
4. The stronger the acid the weaker its conjugate base, the weaker the acid the stronger its conjugate base
5. If nitric acid is a stronger acid than acetic acid, the conjugate base of which is better at creating hydroxide ion?

2. The conjugate acids of bases react with water
1. B^-+ H₂O -> HB + OH^-
2. HB + H₂O -> H₃O⁺ + A^-
3. These equations only describe the behavior of weak bases and their conjugate acids
4. The stronger the base the weaker its conjugate acid, the weaker the base the stronger its conjugate acid
5. If HS^- is a stronger weak base than acetate ion, the conjugate base of which is better at creating hydronium ion?

5. Buffers
1. Maintenance of pH within a narrow range of values is imperative to continued proper performance of many biological functions
2. Buffers are solutions of an acid and its conjugate base (or visa versa) that resist pH changes when either acids or bases are added