Understanding Acids and Bases: Theories and Reactions

1. ACIDS and BASES (unit 11) • Notes start on slide 35 ***

2. ACIDS, BASES & SALTSUnit 11

3. The Arrhenius Theory of Acids and Bases

4. Arrhenius Theory of Acids and Bases: an acid contains hydrogen and ionizes in solutions to produce H+ ions: HCl  H+(aq) + Cl-(aq)

5. Arrhenius Theory of Acids and Bases: a base contains an OH- group and ionizes in solutions to produce OH- ions: NaOH  Na+(aq) + OH-(aq)

6. Neutralization • Neutralization: the combination of H+ with OH- to form water. • H+(aq) + OH-(aq) H2O (l) • Hydrogen ions (H+)in solution form hydronium ions (H3O+)

7. In Reality… H+ + H2O  H3O+ Hydronium Ion (Can be used interchangeably with H+)

8. Commentary on Arrhenius Theory… One problem with the Arrhenius theory is that it’s not comprehensive enough. Some compounds act like acids and bases that don’t fit the standard definition.

9. Bronsted-Lowry Theory of Acids & Bases

10. Bronsted-Lowry Theory of Acids & Bases: • An acid is a proton (H+) donor • A base is a proton (H+) acceptor

11. for example… Proton transfer HCl(aq) + H2O(l) H3O+(aq) + Cl-(aq) Base Acid

12. Water is a proton donor, and thus an acid. another example… CONJUGATE BASE ACID NH3(aq) + H2O(l) NH4+ (aq) + OH- (aq) BASE CONJUGATE ACID Ammonia is a proton acceptor, and thus a base

13. Conjugate acid-base pairs • Conjugate acid-base pairs differ by one proton (H+) A conjugate acid is the particle formed when a base gains a proton. A conjugate base is the particle that remains when an acid gives off a proton.

14. Examples: In the following reactions, label the conjugate acid-base pairs: • H3PO4 + NO2- HNO2 + H2PO4- • CN- + HCO3- HCN + CO32- • HCN + SO32- HSO3- + CN- • H2O + HF  F- + H3O+ acid base c. acid c. base base acid c. acid c. base acid base c. base c. acid c. base c. acid base acid

15. Amphoteric Substances A substance that can act as both an acid and a base (depending on what it is reacting with) is termed amphoteric. Water is a prime example.

16. ACIDS Have a sour taste Change the color of many indicators Are corrosive (react with metals) Neutralize bases Conduct an electric current BASES Have a bitter taste Change the color of many indicators Have a slippery feeling Neutralize acids Conduct an electric current Properties of Acids and Bases

17. Strength of Acids and Bases • A strong acid dissociates completely in sol’n: • HCl  H+(aq) + Cl-(aq) • A weak acid dissociates only partly in sol’n: • HNO2 H+(aq) + NO2-(aq) • A strong base dissociates completely in sol’n: • NaOH  Na+(aq) + OH-(aq) • A weak base dissociates only partly in sol’n: • NH3(aq) + H2O(l) NH4+(aq) + OH-(aq)

18. The Lewis Theory of Acids and Bases

19. The Lewis Theory of Acids & Bases • Lewis acid: a substance that can accept an electron pair to form a covalent bond (electron pair acceptor). • Lewis base: a substance that can donate an electron pair to form a covalent bond (electron pair donor).

20. Neutralization (using Lewis) • Neutralization: the formation of a coordinate covalent bond in which both electrons originated on the same (donor) atom.

21. Example 1: • Ionization of NH3: • NH3 + H2O  NH4+ + OH- base + acid .. .. H H - ..  .. .. + H O H H N H O H H N H + H Acid = electron pair acceptor, base = electron pair donor (to form the covalent bond)

22. Example 2: • Auto-ionization of water: • H2O + H2O  H3O+ + OH- + acid .. .. base .. .. - ..  .. .. + H O H H O H O H H O H + H Acid = electron pair acceptor, base = electron pair donor (to form the covalent bond)

23. Example 3: • Reaction of NH3 with HBr (a Lewis AND a Bronsted-Lowry acid-base reaction): • NH3 + HBr  NH4+ + Br- base + acid .. .. H H -  .. .. .. + H Br H N H Br H N H + H

24. SUMMARY OF ACID-BASE THEORIES

25. Acid-Base Reactions • Neutralization reactions: reactions between acids and metal hydroxide bases which produce a salt and water. • H+ ions and OH- ions combine to form water molecules: • H+(aq) + OH-(aq) H2O(l)

26. Example 1: the reaction of HCl and NaOH (there are 3 ways to write the chemical equation): • Balanced formula unit equation: • HCl(aq) + NaOH(aq) H2O(l) + NaCl(aq) • Total ionic equation: • H+ + Cl- + Na+ + OH- H2O + Na+ + Cl- • Net ionic equation: • H+(aq) + OH-(aq)  H2O(l)

27. Example 2: Write the 3 types of equations for the reaction of hydrobromic acid, HBr, with potassium hydroxide, KOH. • Balanced formula unit equation: • HBr(aq) + KOH(aq) H2O(l) + KBr(aq) • Total ionic equation: • H+ + Br- + K+ + OH- H2O + K+ + Br- • Net ionic equation: • H+(aq) + OH-(aq) H2O(l)

28. Example 3: Write the 3 types of equations for the reaction of nitric acid, HNO3, with calcium hydroxide, Ca(OH)2. • Balanced formula unit equation: • 2HNO3(aq) + Ca(OH)2(aq) 2H2O(l) + Ca(NO3)2(aq) • Total ionic equation: • 2H+ + 2NO3- + Ca2+ + 2OH- 2H2O + Ca2+ + 2NO3- • Net ionic equation: • H+(aq) + OH-(aq) H2O(l)

29. Demos…

30. DEMO: Sponge

31. How does that work?... • The sponge is soaked in Congo red. • Congo red is a dye, a biological stain, and a pH indicator. It has been used as a direct fabric dye for cotton to produce a bright red color. • Scientists use Congo red as a pH indicator (a substance that will change color in the presence of different ion concentrations, [H+])

32. Variety of pH indicators… • There are many different types of pH indicators, such as universal indicator and litmus paper. • Litmus paper comes in red Litmus paper and blue Litmus paper. 

33. Red litmus paper in an acids turns… • Blue litmus paper in a base turns … BLUE RED

34. Demo: tap water vs. dH2O • Both waters have Universal indicator in them (= pH indicator (changes color in the presence of ions), which is a type of weak acids) • The water will change pH, and therefore COLOR (which helps us determine if a solution is acidic or basic) with the addition of HCl (acid) and NaOH (base)

35. Universal Indicator Color Chart pH scale 0 7 14 Acid Neutral Base

36. Why does it take more drops of acid or base to make the tap water change color than it does for the distilled water? • What is distilled water made of? What is tap water made of?

37. Buffered Solutions A solution of a weak acid and a common ion is called a buffered solution.

38. Thus, the solution maintains it’s pH in spite of added acid or base.

39. pH and pOH

40. Ionization of water • Experiments have shown that pure water ionizes very slightly: • 2H2O  H3O+ + OH- • Measurements show that: [H3O+] = [OH-]=1 x 10-7 M • Pure water contains equal concentrations of H3O+ + OH-, so it is neutral.

41. pH • pH is a measure of the concentration of hydronium ions in a solution. • pH = -log [H3O+] or • pH = -log [H+]

42. Example: What is the pH of a solution where [H3O+] = 1 x 10-7 M? • pH = -log [H3O+] • pH = -log(1 x 10-7) • pH = 7

43. Example: What is the pH of a solution where [H3O+] = 1 x 10-5 M? • pH = -log [H3O+] • pH = -log(1 x 10-5) • pH = 5 • When acid is added to water, the [H3O+] increases, and the pH decreases.

44. Example: What is the pH of a solution where [H3O+] = 1 x 10-10 M? • pH = -log [H3O+] • pH = -log(1 x 10-10) • pH = 10 • When base is added to water, the [H3O+] decreases, and the pH increases.

45. The pH Scale 0 7 14 Acid Neutral Base

46. pOH • pOH is a measure of the concentration of hydroxide ions in a solution. • pOH = -log [OH-]

47. Example: What is the pOH of a solution where [OH-] = 1 x 10-5 M? • pOH = -log [OH-] • pOH = -log(1 x 10-5) • pOH = 5

48. How are pH and pOH related? • At every pH, the following relationships hold true: • [H3O+] • [OH-] = 1 x 10-14 M • pH + pOH = 14

49. Example 1: What is the pH of a solution where [H+] = 3.4 x 10-5 M? • pH = -log [H+] • pH = -log(3.4 x 10-5 M) • pH = 4.5

50. Example 2: The pH of a solution is measured to be 8.86. What is the [H+] in this solution? • pH = -log [H+] • 8.86 = -log [H+] • -8.86 = log [H+] • [H+] = antilog (-8.86) • [H+] = 10-8.86 • [H+] = 1.38 x 10-9 M ***you may have to put your calculator into sci mode to get the decimals