Acids and Bases: Introduction

1. Acids and Bases: Introduction Section 19.1

2. Objectives • Identify the physical and chemical properties of acids and bases • Classify solutions as acidic, basic, or neutral • Compare the Arrhenius and Bronsted-Lowry models of acids and bases

3. Key Terms • Acidic solutions • Basic solutions • Arrhenius model • Bronsted-Lowry model • Conjugate acid • Conjugate base • Conjugate acid-base pair • amphoteric

4. Properties of Acids and Bases

5. Self-Ionization • Water molecules can react to form hydronium (H3O+)and hydroxide ions(OH-) • H2O + H2O ↔ H3O+ + OH- • H2O ↔ H+ + OH- • The relative amounts of hydronium ions (also referred to as hydrogen ions, H+) and hydroxide ions determine whether a solution is acid, base, or neutral.

6. Determining acid, base, neutral • If hydrogen ions (H+)and hydroxide ions (OH-) are equal in concentration, the solution is neutral • If there are more hydrogen ions (H+) than hydroxide ions (OH-), the solution is acidic • If there are more hydroxide ions (OH-), than hydrogen (H+) ions the solution is basic

7. Arrhenius model of acid • Arrhenius acid = a substance that contains hydrogen and ionizes to produce hydrogen ions in water. • Example: • HCl (aq) H+ (aq) + Cl-(aq) • HCl + H2O  H3O+ + Cl-

8. Arrhenius model of base • Arrhenius base = a substance that contains a hydroxide group and dissociates to produce a hydroxide ions in aqueous solution. • Example: • NaOH  Na+ + OH-

9. Arrhenius • The Arrhenius acid and base model explains many acids and bases. • However, there are acids and bases that have the appropriate properties, but do not fit the Arrhenius model. • ENTER Brønsted-Lowry . . .

10. Brønsted-Lowry Acid • a hydrogen-ion donor • Example: HCl + H2O →H3O+ + Cl – • HCl is an acid because it donates a hydrogen ion to the water. • Of course, it was an acid according to Arrhenius too. So what is different?

11. Brønsted-Lowry Acid • Example: NH3 + H2O ↔ NH4+ + OH- Water is the acid here, because it donates a hydrogen to the ammonia.

12. Brønsted-Lowry Base • a hydrogen ion acceptor NH3 + H2O ↔ NH4+ + OH- Ammonia is the base here, because it accepts a hydrogen from the water.

13. Brønsted-Lowry Base HCl + H2O  H3O+ + Cl – • Water is the base here because it accepts the hydrogen ion from HCl.

14. Amphoteric • a substance that can act as a Brønsted-Lowry acid OR base. • Example: water as an acid NH3 + H2O ↔ NH4+ + OH- • Example: water as a base HCl + H2O  H3O+ + Cl –

15. Conjugate acid-base pairs • Conjugate acid = species produced when a bases accepts a hydrogen ion from an acid • Conjugate base = species that results from an acid donating a hydrogen ion to a base • NH3 + H2O ↔ NH4+ + OH- • base acid conj a conj b • More on page 599

16. Monoprotic and Polyprotic Acids(all classes) • Monoprotic acid = acid that produces one hydrogen ion (HCl). • Polyprotic acid = acid that produces more than one hydrogen ion. • Diprotic acid = acid that produces 2 hydrogen ions (H2SO4) • Triprotic acid = acid that produces 3 hydrogen ions (H3PO4)

17. Strengths of Acids and Bases Section 19.2

18. Objectives • Relate the strength of an acid or base to its degree of ionization • Compare the strength of a weak acid with the strength of its conjugate base and the strength of a weak base with the strength of its conjugate acid • Explain the relationship between the strengths of acids and bases and the values of their ionization constants

19. Key Terms • Strong acid • Weak acid • Acid ionization constant • Strong base • Weak base • Base ionization constant

20. Strengths of Acids • Strong acid = acid that ionizes completely. • Common examples: HCl, HNO3, H2SO4 Others are listed on table 19-1 HCl  H + + Cl - HCl + H2O  H3O+ + Cl –

21. Weak Acid Weak acid = acid that does not ionizes completely. • Examples: acids other than those memorized as strong. • HCN ↔ H+ + CN - HCN + H2O ↔ H3O++ CN -

22. Strengths of Bases • Strong base = base that dissociates completely to the metal ion and hydroxide ion. • Examples: Group I & II hydroxides • NaOH(s) Na+ + OH-

23. Weak Bases • weak base = base that does NOT dissociates completely to the metal ion and hydroxide ion. • Examples: any base other than group I & II hydroxides NH3 + H2O ↔ NH4+ + OH-

24. Weak vs Strong Electrolyte

25. Concentration • Is concentrated the same as strong? • Is dilute the same as weak?

26. NO • Concentrated is NOT the same as strong when you are discussing acids and bases. • Dilute is NOT the same as weak when you are discussing acids and bases.

27. As you just learned, strong and weak have to do with ionization or dissociation being complete or incomplete. • In an earlier chapter, you learned that concentrated and dilute have to do with the amount of water added, or molarity of a solution.

28. Hydrochloric acid is ALWAYS a strong acid because it ionizes completely. • However, it may be concentrated because little to no water is added. • Or- it may be dilute if a lot of water is added. • HCl can be a concentrated strong acid or a dilute strong acid.

29. Ammonia is ALWAYS a weak base because it doesn’t dissociate completely. • However, it may be concentrated because little to no water is added. • Or- it may be dilute if a lot of water is added. • Ammonia can be a concentrated weak base or a dilute weak base.

30. Acid ionization constant, Ka • Just like Keq • HCN (aq) + H2O (l) ↔ H3O+(aq)+ CN-(aq) • Ka = [H3O] [CN] [HCN] • Mathematically, a large Ka means larger numerator more ions; • so the larger the Ka , the more the acid ionizes and the stronger it is.

31. Acid ionization constant, Ka • Table 19-2 page 605 • Polyprotic acids have a Ka for each H

32. Base ionization constant, Kb • Just like Ka • CH3NH3(aq) + H2O (l) ↔ CH3NH4+(aq)+ OH-(aq) • Kb = [CH3NH4+][OH-] [CH3NH3] • Mathematically, a large Kb means larger numerator more ions; • so the larger the Kb , the more the base ionizes and the stronger it is.