Chapter 19 Acids and Bases

Lecture
19.1 Properties of Acids and Bases

1. Have sour tastes.

2. Acetic acid (ethanoic acid) in vinegar. Citric acid in fruits. Aqueous solutions conduct electricity, some strong, some weak. Acids cause indicators to change color. Metals dissolve in acids to form hydrogen gas. Acids react with hydroxide ions to form water and a salt.

3. Bases react with acids to form water and salt. Bases neutralize acids. Have bitter taste, are slippery, and can be strong or weak electrolytes.

19.2 Names and Formulas of Acids and Bases.

1. Acid: a compound that produces hydrogen when dissolved in water

2. General formula HX. X = mon or poly atomic anion.

3. HCl (g) dissolves in water to form HCl (aq). HNO₃ is always aqueous.

Rules for naming acids

• Anion -ide. Begin with hydro and add -ic. Eg., HCl - hydrochloric acid.
• Anion -ite. Use anion stem with -ous acid. Eg., HNO₂ - Nitrous acid.
• Anion -ate. Use anion stem with -ic acid. HNO₃ - Nitric acid.

5. Naming can work in reverse. Write formulas given names. Eg., Chloric acid. HClO₃. Hdrobromic acid. HBr, etc.

6. Base: a compound that produces hydroxide ions when dissolved in water.

7. Bases are named like other ionic compounds. Cation followed by name of anion. Eg., NaOH or Ca(OH)₂

19.3 Hydrogen Ions from Water

1. Water occasionally dissociates:

H₂O <======>H⁺ + OH^-

2. Reaction gives rise to hydroxide and hydrogen ion called self ionization of water.

3. Hydrogen ions are always joined to water molecules as H₃O⁺. (Hydronium ion).
H20 + H2O ---> H₃O^{+  +} OH^-
4. In pure water [H⁺] = [OH^-] which = 1x10^-7 mol/L. Any aqueous solution in which [H] = [OH] is a neutral solution.

5. H⁺ + OH^- <====> H₂0

6. According to Le Chatelier's principle, if increase one, decrease other and more water forms in process.

7. Product of [H]x[OH] = 1.0 x 10^-14 (mol/L)^{2 = Kw}

8. This is called the ion-product constant for water. Kw.

9. Not all solutions are neutral. When dissolved in water they add either H+ or OH-. Eg., HCl or NaOH.

10. [H]>[OH] = acid. [OH]>[H] = basic solution or alkaline solution.

11. Example 2.

19.4 The pH Concept

1. Scale for expressing [H+]

2. 0 is strongly acidic. 14 is strongly basic.

3. The pH of a solution is the negative logarithm of the hydrogen-ion concentration.

4. pH = -log[H+]

5. Neutral solution[H+]=1.0x10^-7

pH=-log (1.0x10^-7)=-(log 1 + log 10^-7) =-(0.0 + (-7))=7.0

6. pOH = -log of [OH-]

7. pH + pOH = 14

8. Example 3 and 4

9. See page 543 for solution pH.

10. See table 19.2 page 543.
19.5 Calculating pH Values - Honors
1. See examples 5 & 6.

19.6 Arrhenius Acids and Bases.

1. 1887 Svante Arrhenius said: acids are compounds containing hydrogen that ionize to yield hydrogen ions in aqueous solution. Bases are compounds that ionize to yield hydroxide ions in aqueous solution.

2. 1 hydrogen is monoprotic - Nictric, HCl

3. 2 hydrogens is diprotic - Sulfuric acid.

4. 3 hydrogens is triprotic - Phosphoric acid

5. Not all H's in compounds are ionizable. Only hydrogens in very polar bonds are ionizable - not all hydrogens.

7. Methane is not polar and does not ionizes. H's do not dissolve.

8. Ethanoic acid (acetic) has a polar H which dissolves. See page 548.

9. Table 19.4 shows common bases. (Bitter and slippery).

10. Eg., sodium hydroxide formed by dissolving Na metal in water. K also.

11. Both are 1A elements - Alkali metals and are VERY soluble in water.

12. Calcium hydroxide and Magnesium hydroxide (milk of magnesia) are 2A elements - alkaline earth metals. They are slightly soluble. Do 15 & 16.

19.7 Bronsted-Lowry Acids and Bases.
(Arrhenius does not explain all situations, eg., Ammonia forms base without OH)

1. The Bronsted-Lowry Theoy defines and acid as a hydrogen -ion donor. Bases are hydrogen-ion acceptors.

2. This covers a wider range of acids and bases than Arrhenius.

3. NH₃^(base) + H₂0(acid)<=====> NH₄^{+(conjugate acid)} + OH^{-(conjugate base)}

Ammonia-base, water=acid.

Ammonium =conjugate acid, OH-is conjugate base. (when reversed-).

4. Ammonia is a gas which is very soluble in water. H+ ions are transferred from water to the ammonia. OH- concentration is greater than in pure water, therefore becomes a base.

5. If heat then drive ammonia out and force reaction from right to left. Ammonium ion then acts like an acid, giving up the hydrogen. Water, then, acts like a base, accepting the hydrogen.

6. Conjgate acid-base pair is 2 sbstances that are related by the loss or gain of a single hydrogen ion. See page 551.

7. HCl + H₂0 <======> H₃O+ + Cl^-
 

8. A substance which can be either an acid or base is amphoteric. Water. With HCl water is a base (accepts proton). With NH₃ water is an acid (donates proton).
19.8 Measuring pH
1. For preliminary small volume samples indicators are used.
2. For precise measurements a pH meter is used.
3. An indicator is a weak acid or base that undergoes dissociation in a know pH range. The acid or base is a different color from its conjugate acid or base.
4. Each indicator is used over a range of 2 pH units. (see page 553).
5. Indicator strips, impregnated with the indicators, may be used instead.
6. A pH meter uses electrodes with a constant voltage. The H3O+ concentration changes the mv and the meter measures the difference between the electrodes as pH. Is accurate within .01 pH units.
19.9 Lewis Acids and Bases
1. Gilber Lewis 875-1946: 3rd theory
2. A Lewis acid is a substance that can accept a a pair of electrons to form a convalent bond.  A Lewis base is a substance that can donate a pair of electrons to form a covalent bond.

3. See Example 7
19.10 Strengths of Acids and Bases
1. Acids are classified as strong or weak depending on degree to which ionize in water.
2. Strong acids are completely ionized in aqueous solutions. Eg., HCl or H2SO4, page 557.
3. Weak acids ionize only slightly in aqueous solutions.Eg., ethanoic acid, age 557. (1% molecules ionized).
4. See table 19.5, page 557 for examples.
5. The base dissociation constant, Kb, is the ratio of concentrations: [conjugate acid]x[OH-]/[conjgate base]

6. Also called ionization constant.
7. If number is small, then amount of dissciation is small and acid is weak.
8. See table 19.6, page 558
9. Diprotic and triprotic acids lose protons one at a time. Each has separate dissciation constant.
10. Strong bases dissociate completely into metal ions and hydroxide ions in aqueous solution.
11. Weak bases react with water to form the conjugate acid ofthe base and the hydroxide ion:

12. The base dissociation constant, Kb, is the ratio of concentrations: [conjugate acid]x[OH-]/[conjugatebase]

13. Weak bases have small dissociation constants.
14. Concentrated and dilute tell how much of an acid or base dissolves.
15. Strong or weak tell how much ionization occurs.
16. Acids may be strong, such as HCl, but dilute, or weak, sucha as ethanoic acid and concentrated.
17. Bases may be strong, Ca(OH)2 or Mg(OH)2, but dilute because the do not dissolve very well.
18. Combinations result in different effects of acits or bases.
19. Ammonia is a weak base, whether concentrated or dilut.
19.11 Calculating Dissociation Constants
1. We can calculate Ka with experimental data with 2 conditions.