1. Acids and Bases
Chapter 14 and Section 15.2

2. Objectives Identify key terms and concepts.
Identify an acid and a base using Bronstead- Lowery Acid/Base Theory.
Identify a Lewis acid and base.
Explain the acid-base properties of water.
Name acids and bases.
Explain the properties of acids, bases, and salts.
Explain how a buffer works.
Identify acid-base conjugate pairs.
Determine the products of a neutralization reaction.

3. Objectives Calculate the pH and hydronium concentration of a solution.
Calculate acid/base ionization constants.
Determine the percent ionization for an acid or a base.
Explain how a salt will affect the pH of a solution.
Identify the acid-base properties of an oxide and hydroxide.
Determine the relative strength of an acid-base conjugate pair.

4. Objectives
Explain how a buffer works and how to prepare a buffer system for a specific pH.
Calculate the pH of a buffer solution.
Calculate the equivalence point for acid-base titrations.

5. Acids Cause litmus paper to turn red Tastes sour
Dissolves metals like Zn and Fe to produce hydrogen gas

6. Acids
For an acid containing a hydrogen and an anion: use the prefix hydro and change the –ide of the nonmetal to –ic and add the word acid.
HCl = Hydrogen Chloride is Hydrochloric Acid
For an acid containing a polyatomic ion, name the anion after the polyatomic ion, changing the –ate to –ic or the –ite to –ous and add the word acid.
HNO3 = Hydrogen Nitrate is Nitric Acid
HNO2 = Hydrogen Nitrite is Nitrous Acid

7. Write the Name or Formula for each of the Following Acids
H2SO4 HF
HBr
H3PO4
Acetic acid
Carbonic Acid
Hydroiodic Acid
Phosphorous Acid

8. Acids Strong Acids Weak Acids
Strong electrolytes that will completely ionize in water
H2SO4  H+ + HSO4-
Only a few strong acids
HCl, HNO3, H2SO4, HBr, HI, HClO4
Weak Acids
Will only partially ionize in water (weak electrolytes)
Most acids are weak
HF, acetic acid, citric acid, salicylic acid, oxalic acid
HF + H2O ↔ F- + H3O+

9. Bases Cause litmus paper to turn blue Tastes bitter
Name like a normal compound
Ca(OH)2  Calcium Hydroxide
NaOH  Sodium Hydroxide

10. Bases Strong Bases Weak Bases
Strong electrolytes that will completely ionize in water
NaOH  Na+ + OH-
LiOH, NaOH, KOH, Ca(OH)2, Sr(OH)2, Ba(OH)2
Weak Bases
Weak Bases do not furnish OH- ions by dissociation. Instead, they react with water to generate OH- ions.
Weak electrolytes
NH3 + H2O ↔ NH4 + OH-

11. Acids
Bases

12. Acid-Base Indicators pH meters Litmus and pH paper Universal Indicator
Phenolphthalein
Methyl Red
Thymol Blue

13. Percent Ionization Percent Ionization
The higher the percent ionization, the stronger the acid
Ionized acid concentration at equilibrium
Initial concentration of acid
% Ionization = x 100

14. The pH of a 0. 30M solution of acetic acid is 2. 79
The pH of a 0.30M solution of acetic acid is What is the percent ionization?

15. The pH of a 0. 15M solution of hydrofluoric acid is 3. 25
The pH of a 0.15M solution of hydrofluoric acid is What is the percent ionization?

16. The Arrhenius Theory Acid
A substance that ionizes into hydrogen ions and anions in water
HCl  H+ + Cl-
Base
A substance that releases hydroxide ions in water
NaOH  Na+ + OH-
Stated the reaction of an acid and a base will produce salt and water
HCl + NaOH  NaCl + H2O

17. Problems with the Arrhenius Theory
A hydrogen ion does not exist in water. Instead, the H+ ion attaches to one of the oxygen loan pairs of electrons on water to make a hydronium ion.
Does not explain the basicity of ammonia of where the hydroxide ion comes from when ammonia is dissolved in an aqueous solution.
NH3 + H2O  NH4 + OH-
Applies only to reactions in aqueous solution (water).

18. Bronstead-Lowry Acid-Base Theory
Proton Donor
HCl + H2O  H3O+ + Cl-
Base
Proton Acceptor
NH3 + H2O  NH4 + OH-

19. Identify the acid and base in the following reactions using Bronstead-Lowery Acid/Base Theory
S2- + H2O  OH- + HS-
PO43- + HNO3  NO3- + HPO42-
C2O42- + HC2H3O2  HC2O4- + C2H3O2-
HCO2H + OH-  CO2H- + H2O
HCl + H2O  H3O + Cl-

20. Lewis Acids and Bases Lewis Acid Lewis Base
Accepts a pair of electrons
Lewis Base
Donates a pair of electrons
Much more general description of acids and bases
The reaction does not have to involve the transfer of a hydrogen.

21. Identify the acid and base in the following reactions using Lewis acid/base Theory
Cr3+ + 6H2O  Cr(H2O)63+ BeCl + H2O  BeCl(H2O)2 AlCl3 + CH3Cl  AlCl3CH3Cl SO2 + H2O SO3H2 BF3 + H2O  BF3H2O

22. What are your questions?

23. The pH scale Measures how basic or acidic a solution is
Measures the concentration of hydronium present
pH = -log[H3O+]
pOH = -log[OH-]
pH + pOH = 14
pH > 7, basic solution
pH < 7, acidic solution
pH = 7, neutral solution

25. What is the pH of the solutions containing the following concentrations of hydronium ions? What is the pOH of the solutions?
1x10-2 M 3.5x10-3 M 4x10-7 M 2x10-6 M 1x10-9 M

26. What is the hydronium and hydroxide ion concentrations for the solutions with the following pH?

27. Reactions of Acids and Bases
Neutralization
The reaction between a strong acid and a strong base with produces salt and water
Double replacement reaction
Salt
An ionic compound that does not have H+ or OH- ions

28. Complete the following neutralization reactions
Ca(OH)2 + HNO3 
NH3 + H2SO4 
Mg(OH)2 + H3PO4 
NaOH + HC2H3O4 

29. Neutralization Reactions – Gas Forming
If a gas forms, such as CO2 (from H2CO3), SO2 (from H2SO3), or NH3 (from NH4OH), it is called a gas forming reaction.
NH4Br + KOH 
HC2H3O2 + NaHCO3 
HCl + LiHSO3 

30. Molecular Structure and the Strength of Acids
Properties affecting the strength of acids
Properties of the solvent
Temperature
Molecular structure of the binary acid
Ionization
The stronger the bond, the weaker the acid
Going down a column on the periodic table, the size of the molecule increases, the bond strength decreases, and the strength of the acid increases.
Higher the polarity of the molecule, the stronger the acid
Going across a row of the periodic table, the electronegativity increases, the bond polarity increases, and the acid strength increases.

31. Molecular Structure and the Strength of Acids
Hydrohalic Acids
Acids containing a halogen
The bond strength in these acids has a larger influence over the strength of the acid than the polarity of the molecule
HF < HCl < HBr < HI

32. Molecular Structure and the Strength of Acids
Oxoacids
Contain H, O, and a third element
Oxoacids with different central atoms from the same group on the periodic table with the same oxidation number
As the electronegativity of the central atom increases, the acid strength increases
Oxoacids with the same central atom but different numbers of attached groups
As the oxidation number of the central atom increases, so does the acid strength
HClO4 > HClO3 > HClO2 > HClO

33. Which of the following acids has the greatest strength?
NH3 or PH3
HI or H2Te
HSO3- or H2SO3
H3AsO3 or H3AsO4
HSO4- or HSeO4-

34. Conjugate Acids and Bases
Conjugate Acid-Base Pair
Pair of molecules or ions related by the loss of gain of a H+

35. Conjugate Acids and Bases

36. Conjugate Acids and Bases
Properties of conjugate acid-base pairs
If an acid is strong, its conjugate base is very weak.
H3O+ is the strongest acid that can exist in water. Acids stronger than H3O+ will react completely with water to form H3O+ and a conjugate base. Acids weaker than H3O+ will react with water to a lesser extent to produce H3O+ and a conjugate acid (reversible reaction).
OH- is the strongest base that can exist in water. Bases stronger than OH- will react completely with water to produce OH- and a conjugate acid.

38. Identify the conjugate acid-base pairs in the following reactions
NH3 + H2O  NH4+ + OH- HCl + H2O  H3O + Cl- H2O + CH3CO2H  H3O + CH3CO2- H2O + H2SO4  H3O+ + HSO4-

39. Identify the conjugate acid-base pairs in the following reactions
S2- + H2O  OH- + HS- PO43- + HNO3  NO3- + HPO42- C2O42- + HC2H3O2  HC2O4- + C2H3O2- HCO2H + OH-  CO2H- + H2O

40. What are your questions?

41. Acid-Base Properties of Salts
Strong electrolytes that dissolve completely in water
NaCl, NaNO3, KBr
Salt Hydrolysis
The reaction of an anion or a cation of a salt, or both, with water.
Affects the pH of water

42. Acid-Base Properties of Salts
Salts that produce neutral solutions
CaCl2(aq) Ca+(aq) + 2Cl-(aq)
Salts that produce basic solutions
CH3COONa(aq) Na+(aq) + CH3COO-(aq)
CH3COO- (aq) + H2O(l) OH-(aq) + CH3COOH(aq)
Salts that produce acidic solutions
NH4Cl(aq) NH4+(aq) + Cl-(aq)
NH4+(aq) + H2O(l) NH3 (aq) + H+(aq)

43. Acid-Based Properties of Salts
Kb > Ka, then the solution is basic because the anion will hydrolyze to a great extent than the cation
Kb < Ka, then the solution will be acidic because cation hydrolysis will be greater than anion hydrolysis
Kb = Ka, then the solution will be near neutral pH

44. Predict weather a salt solution is acids, basic, or neutral?
NaCN
LiF
KNO3
NaHCO3
NaC7H5O2

45. Acid-Base Properties of Oxides
Basic oxides will form bases when reacted with water.
Na2O + H2O  NaOH
Acidic oxides will form an acid when reacted with water.
SO3 + H2O  H2SO4
Reactions between acidic/basic oxides and acids/bases produce salt and water.
Na2O + 2HCl  2NaCl + H2O
SO3 + 2LiOH  Li2SO4 + H2O
Amphoteric oxides can act as both an acid and a base.

46. Acidic, Basic, and Amphoteric Oxides

47. Acid-Base Properties of Hydroxides
Alkali and alkaline metal hydroxides are basic – Except for Be(OH)2
Many metal hydroxides are amphoteric
Al(OH)3, Sn(OH)2, Pb(OH)2, Cr(OH)2, Cu(OH)2, Zn(OH)2, Cd(OH)2
Sn(OH)2 + 2H+  Sn2+ + H2O
Sn(OH)2 + 2OH-  Sn(OH)42-
All amphoteric hydroxides are insoluble.

48. What are your questions?

49. Acid Base Properties of Water
Can act as either an acid or a base.
Base
H2O(l) + CH3CO2H (aq) H3O+(aq) + CH3CO2- (aq)
Acid
NH3 (aq) + H2O(l) NH4+ (aq) + OH- (aq)
Water can undergo ionization to a small extent.
H2O(l) H+(aq) + OH-(aq)
During this ionization, water can act as both the acid and the base.
H2O(l) + H2O(l) H3O+(aq) + OH-(aq)

50. Ionization Constants Acid Ionization Constant
Equilibrium constant for the ionization of an acid
H2O(l) + HA(aq) H3O+(aq) + A-(aq)
Base Ionization Constant
Equilibrium constant for the ionization of a base
NH3 (aq) + H2O(l) NH4+ (aq) + OH- (aq)
[H3O+][A-]
[HA]
Ka =
[NH4+ ][OH-]
[NH3]
Kb =

52. Ionization Constants
For conjugate acid-base pairs, the ionization constants are related using the following equation
CH3COOH (aq) H+(aq) + CH3COO- (aq)
CH3COO- (aq) + H2O(l) CH3COOH (aq) + OH- (aq)
H2O(l) H+(aq) + OH- (aq)
Kw = KaKb

53. HC2H3O2(aq) C2H3O2-(aq) + H+(aq)
Calculate the pH of a 0.5M solution of acetic acid in aqueous solution:
HC2H3O2(aq) C2H3O2-(aq) + H+(aq)

54. Calculate the pH of a 0.65M solution of HF in aqueous solution:
HF(aq) F-(aq) + H+(aq)

55. What is the pH of a 0.25M solution of ammonia?

56. The pH of a 0. 30M solution of acetic acid is 2. 79
The pH of a 0.30M solution of acetic acid is Calculate the Ka of the acid?

57. The pH of a 0. 15M solution of hydrofluoric acid is 3. 25
The pH of a 0.15M solution of hydrofluoric acid is Calculate the Ka of the acid?

58. The pH of a 0. 75M solution of methylamine, CH3NH2, is 8. 78
The pH of a 0.75M solution of methylamine, CH3NH2, is Calculate the Kb of the base?

59. Diprotic and polyprotic acids
Produce more than 1 hydrogen per mole when ionized.
Ionization occurs in several steps.
H2SO3(aq) H+(aq) + HSO3-(aq)
HSO3-(aq) H+(aq) + SO32-(aq)

60. When sulfurous acid is dissolved in water, it dissociates in two steps
When sulfurous acid is dissolved in water, it dissociates in two steps. Calculate the concentrations of all species present at equilibrium in a 0.25M solution. H2SO3(aq) H+(aq) + HSO3-(aq) HSO3-(aq) H+(aq) + SO32-(aq)

61. What are your questions?

62. Conjugate Acids and Bases

63. Buffers
Buffer solutions resist changes in pH when small amounts of acids or bases are added
Buffers are made of a weak acid or base and it’s salt (acid-base conjugate pair)
Buffering Capacity

64. Buffers

65. Identify the conjugate acid and/or base for the following chemical
Identify the conjugate acid and/or base for the following chemical. NH3 HCN C2H3O2- H2PO3- HCO3-

66. Instructions for making up a buffer say to mix 60mL of 0
Instructions for making up a buffer say to mix 60mL of 0.1M NH3 with 40mL of 0.1M NH4Cl. What is the pH of the buffer?

67. Instructions for making up a buffer say to mix 100mL of 0
Instructions for making up a buffer say to mix 100mL of 0.25M HC2H3O2 with 75mL of 0.3M NaC2H3O2. What is the pH of the buffer?

68. Buffers [base] pH = pKa + log [acid] pKa = -log Ka
Henderson-Hasselbalch Equation
Relates pH of a buffer for different concentrations of the conjugate acid-base
If the molar concentration of the acid ≈ base, then the pH ≈ pKa
[base]
[acid]
pH = pKa + log
pKa = -log Ka

70. What must the ratio of the concentration of the HCO3- and H2CO3 buffer system found in blood be for the pH to be 7.4?

71. What acid-base conjugate pair could be used to make a buffer with a pH of 3.7? How would you prepare the buffer solution? H3PO4 ↔ H2PO4- + H3O+ Ka = 7.5x10-3 H2S ↔ HS- + H3O+ Ka = 1x10-7 HCO2H ↔ HCO2- + H3O+ Ka = 1.8x10-4

72. Calculate the pH of a buffer solution using ammonia and ammonium to which 3mL of 0.10M HNO3 is added.

73. Calculate the pH of a buffer solution using acetic acid/sodium acetate to which 9.5mL of 0.10M HCl is added.

74. Acid-Base Titrations Titration
Addition of a solution of a known concentration (standard solution) to a solution of unknown concentration until the reaction is complete (reaches its equivalence point)

76. Acid-Base Titrations Strong Acid-Base Titrations
Weak Acid-Base Titrations
Strong Acid-Weak Base Titration
Weak Acid-Strong Base

77. Calculate the pH of a solution in which 10mL of 0
Calculate the pH of a solution in which 10mL of 0.1M NaOH is added to 25.0mL of 0.1M HCl.

78. Calculate the pH of a solution when 25mL of 0
Calculate the pH of a solution when 25mL of 0.1M nicotinic acid (HC5H4NO2) is titrated by 20mL of 0.1M NaOH. The Ka for nicotinic acid is 1.4x10-5.

79. What is the pH of a solution when 35mL of 0
What is the pH of a solution when 35mL of 0.2M ammonia is titrated by28mL 0.12M HCl. The Kb for ammonia is 1.8x10-5.

80. What is the pH at the equivalence point when 25mL of 0
What is the pH at the equivalence point when 25mL of 0.1M HF is titrated by 0.15M NaOH?

81. What are your questions?