1. Chapter 14 Acids and Bases

2. Chapter 14 Section 1 – Properties of Acids and Bases
Section 2 – Acid Base Theories
Section 3 – Acid Base Reactions

3. 14.1 Properties of Acids and Bases
List five general properties of aqueous acids and bases.
Name common binary acids and oxyacids, given their chemical formulas.
List five acids commonly used in industry and the laboratory, and give two properties of each.
Define acid and base according to Arrhenius’s theory of ionization.
Explain the differences between strong and weak acids and bases.

4. Properties of: Acids Bases
Sour taste
Conducts electricity
Turns litmus paper red
Reacts with bases to produce salts and water
Reacts with some metals and releases hydrogen gas
Can you think of a reaction that this occurs?
Bitter taste
Feels slippery
Conducts electric current
Turns litmus paper blue
Reacts with acids to produce salts and water

5. Binary Acids Contains only two different elements Nomenclature:
Hydrogen and an electronegative element (usually a halogen)
Nomenclature:
hydro - _________ - ic acid

6. Diatomic Nomenclature

7. Oxyacid
Contains hydrogen, oxygen, and a third element (hydrogen with a polyatomic ion)
Nomenclature:

8. Acid Names

9. Oxyacids

10. Common Industrial Acids
Sulfuric Acid
Sulfuric acid is the most commonly produced industrial chemical in the world.
Nitric Acid
Phosphoric Acid
Hydrochloric Acid
Concentrated solutions of hydrochloric acid are commonly referred to as muriatic acid.
Acetic Acid
Pure acetic acid is a clear, colorless, and pungent-smelling liquid known as glacial acetic acid.

11. Arrhenius Acids and Bases
Increases concentration of H+ ions in solution
Arrhenius Bases:
Increases concentration of OH- ions in solution

12. Arrhenius Acid Base Video

13. Arrhenius Acids Molecular compounds with ionizable hydrogen atoms
Water solutions are known as aqueous acids
Electrolytes

14. Acid Strength Strong acid: Weak acid:
Ionizes completely in solution and is an electrolyte
Example: HCl, HClO4, HNO3
Weak acid:
Releases few hydrogen ions in solution
Hydronium ions, anions and dissolved acid molecules present
Examples: HCN, Organic acids – HC2H3O2

15. Aqueous Acids

16. Base Strength Strong bases: Weak bases:
Ionic compounds containing metal cation and hydroxide ion (OH-)
Dissociates in water
Weak bases:
Molecular compounds do not follow Arrhenius definition:
Ammonia (NH3)
Produces hydroxide ions when it reacts with water molecules

17. Base Strength

18. Acid Base Strength Video

19. Acidic solution has greater [H3O+] Basic solution has greater [OH–]

20. 14.2 Acid Base Theories
Define and recognize Brønsted-Lowry acids and bases.
Define a Lewis acid and a Lewis base.
Name compounds that are acids under the Lewis definition but are not acids under the Brønsted-Lowry definition.

21. Bronsted-Lowry Acid Bronsted-Lowry Acid: Proton (H+) donor
Hydrogen chloride acts as a Bronsted-Lowry acid when it reacts with ammonia.
Water can also act as a Bronsted-Lowry acid

22. Bronsted-Lowry Base Bronsted-Lowry Base: Proton acceptor
Ammonia accepts a proton from hydrochloric acid.
Hydroxide ions produced in solution act as a Bronsted- Lowry base

23. Bronsted-Lowry Acid Base Reactions
Protons are transferred from one reactant (the acid) to another (the base)
acid base

24. Brønsted-Lowry Acid Base Video

25. Monoprotic Acids
Can donate only one proton (hydrogen ion) per molecule
One ionization step

26. Monoprotic and Diprotic Acids

27. Polyprotic Acids Donates more than one proton per molecules
Multiple ionization steps
Diprotic – donates 2 protons Ex:
Triprotic – donates 3 protons Ex:
Sulfuric acid solutions contain H3O+, HSO4-, SO4- ions 1. 2.

28. Lewis Acid Lewis acid: Lewis base:
Atom, ion, or molecule that ACCEPTS an ELECTRON PAIR to form a covalent bond
A proton (hydrogen ion) is a Lewis acid
Lewis base:
Atom, ion, or molecule that DONATES an ELECTRON PAIR to form a covalent bond

29. Lewis Acid A lewis acid might not include hydrogen
Silver as a lewis acid:

30. Lewis Acid Base Video

31. Acid and Base Definitions

32. Acid Base Definitions Video

33. 14.3 Acid Base Reactions
Describe a conjugate acid, a conjugate base, and an amphoteric compound.
Explain the process of neutralization.
Define acid rain, give examples of compounds that can cause acid rain, and describe effects of acid rain.

34. Conjugate Acid – Base acid conjugate base Conjugate Base:
The species that remains after a Bronsted-Lowry acid has given up a proton
Conjugate Acid:
The species that remains after a Bronsted-Lowry base has accepted a proton
acid conjugate
base

35. Conjugate Acid Base Pairs
Match up the acid-base pairs (proton donor-acceptor pairs)
acid1 base base acid2

36. Strength of Acid Base Pairs
The stronger the acid, the weaker the conjugate base
The stronger the base, the weaker the conjugate acid
strong acid base acid weak base

37. Proton transfer favors the production of the weaker acid and base.
stronger acid stronger base weaker acid weaker base
weaker acid weaker base stronger acid stronger base

38. Acid Base Strength

39. Amphoteric Any species that can react as either an acid or base
Example: water
acid base acid base1
base acid acid base2

40. Amphoteric Water Video

41. Amphoteric Compounds
Covalently bonded –OH group in an acid is referred to as a hydroxyl group
Molecular compounds with hydroxyl groups can be acidic or amphoteric
The behavior of the compound is affected by the number of oxygen atoms bonded to the atom connected to the –OH group

42. Oxyacids of Chlorine

43. Neutralization Reactions
What does it mean to neutralize something?
Neutralization reactions:
Hydronium and hydroxide ions react to form water
The left over cation and anion in solution produce a salt (ionic compound)

44. Neutralization Reactions

45. Neutralization Reaction Video

46. Acid Rain
NO, NO2, CO2, SO2, and SO3 gases from industrial processes can dissolve in atmospheric water to produce acidic solutions.
Very acidic rain is known as acid rain.
Acid rain can erode statues and affect ecosystems.

47. Chapter 15 Acid Base Titration and pH

48. Chapter 15 Section 1 – Aqueous Solutions and the Concept of pH
Section 2 – Determining pH and Titrations

49. 15.1 Aqueous Solutions and pH
Describe the self-ionization of water.
Define pH, and give the pH of a neutral solution at 25°C.
Explain and use the pH scale.
Given [H3O+] or [OH−], find pH.
Given pH, find [H3O+] or [OH−].

50. Self Ionization of Water
Two water molecules produce a hydronium ion and hydroxide ion by proton transfer
In water at 25°C, [H3O+] = 1.0 ×10−7 M and [OH−] = 1.0 × 10−7 M
The ionization constant of water, Kw
Kw = [H3O+][OH−]

51. Kw = [H3O+][OH−] = (1.0 × 10−7)(1.0 × 10−7) = 1.0 × 10−14
At 25OC
Kw = [H3O+][OH−] = (1.0 × 10−7)(1.0 × 10−7) = 1.0 × 10−14
Kw = 1.0 x 10-14
Kw increases as temperature increases

52. Ion Concentration neutral acidic basic [H3O+] = [OH−]
[H3O+] > 1.0 × 10−7 M
[OH−] > [H3O+]
basic
[OH−] > 1.0 × 10−7 M

53. Calculating Concentration
Strong acids and bases are considered completely ionized or dissociated in weak aqueous solutions.
1 mol mol mol
1.0 × 10−2 M NaOH => [OH−] = 1.0 × 10−2 M
[H3O+] is calculated using Kw
Kw = [H3O+][OH−] = 1.0 × 10−14

54. Example Problem 1 [H3O+] = ______________ Unknown: [OH-] = ?
Given: [HCl] = 2.0 × 10−4 M
[H3O+] = ______________
Unknown: [OH-] = ?
Kw = [H3O+][OH−] = 1.0 × 10−14

55. Example Problem 2
A 1.0  10–4 M solution of nitric acid has been prepared for a laboratory experiment.
Calculate [H3O+]
Calculate [OH–]

56. Example Problem 2 Solution
Given:
Concentration of the solution = 1.0 × 10−4 M HNO3
Unknown: [H3O+]
[OH−]
Solution:
Assume HNO3 dissociates completely
1 mol mol mol mol

57. Example Problem 2 Solution
Remember: [H3O+][OH−] = 1.0 × 10−14

58. pH = −log [H3O+] = −log(1 × 10−7) = −(−7.0) = 7.0
The pH of a solution is defined as the negative of the common logarithm of the hydronium ion concentration, [H3O+].
pH = −log [H3O+]
Example:
a neutral solution has a [H3O+] = 1×10−7
pH = −log [H3O+] = −log(1 × 10−7) = −(−7.0) = 7.0

59. pH Values as Specified [H3O+]

60. pOH pOH = −log [OH–] pH + pOH = 14.0
The pOH of a solution is defined as the negative of the common logarithm of the hydroxide ion concentration, [OH−].
pOH = −log [OH–]
pH + pOH = 14.0
Example:
a neutral solution has a [OH–] = 1×10−7
the pH of this solution is?

61. pOH Video

62. The pH Scale

63. Approximate pH Range of Common Materials

64. Comparing pH and pOH Video

65. Significant Figures
There must be as many significant figures to the right of the decimal as there are in the number whose logarithm was found.
Example: [H3O+] = 1 × 10−7
one significant figure
pH = 7.0

66. The Circle of pH pH pOH [ H3O+] [ OH-] -log [H3O+] antilog (-pH)
antilog (-pOH)
-log [OH-]
= 1.0x10-14
+ pOH
= 14

67. pH of Weak Acids and Bases
The pH of solutions of weak acids and weak bases must be measured experimentally.
The [H3O+] and [OH−] can then be calculated from the measured pH values.

68. pH Values of Some Common Materials

69. 15.2 Determining pH and Titrations
Describe how an acid-base indicator functions.
Explain how to carry out an acid-base titration.
Calculate the molarity of a solution from titration data.

70. Indicators
Acid-base indicators: compounds whose colors are sensitive to pH.
The pH range over which an indicator changes color is called its transition interval.

71. pH Meters
pH meter determines the pH of a solution by measuring the voltage between the two electrodes that are placed in the solution.
The voltage changes as the hydronium ion concentration in the solution changes.
Measures pH more precisely than indicators

72. Color Ranges of Indicators

73. Color Ranges of Indicators

74. Color Ranges of Indicators

75. Antacids Video with Methyl Orange

76. H3O+(aq) + OH−(aq) 2H2O(l)
Titration
Neutralization occurs when hydronium ions and hydroxide ions are supplied in equal numbers by reactants.
H3O+(aq) + OH−(aq) 2H2O(l)
Titration: the controlled addition and measurement of the amount of a solution of known concentration required to react completely with a measured amount of a solution of unknown concentration.

77. Titration Points
equivalence point: point at which the two solutions used in a titration are present in chemically equivalent amounts
end point: point in a titration at which an indicator changes color

78. Which indicator do I choose?
pH less than 7
Indicators that change color at pH lower than 7 are used to determine the equivalence point of strong-acid/weak-base titrations.
strong-acid/weak-base titration = acidic.
pH at 7
Indicators that undergo transition at about pH 7 are used to determine the equivalence point of strong-acid/strong base titrations.
strong acids/strong bases = salt solution with a pH of 7.

79. Which indicator do I choose?
pH greater than 7
Indicators that change color at pH higher than 7 are used to determine the equivalence point of weak-acid/strong-base titrations.
weak-acid/strong-base = basic

80. Titration Curve Strong Acid and a Strong Base
Equivalence Point:
pH at 7

81. Titration Curve Weak Acid and a Strong Base
Equivalence Point:
pH higher than 7

82. Titration Curve Strong Acid and a Weak Base
Equivalence Point:
pH less than 7

83. Molarity and Titration
standard solution: solution that contains the precisely known concentration of a solute
primary standard: highly purified solid compound used to check the concentration of the known solution
The standard solution can be used to determine the molarity of another solution by titration.

84. Performing a Titration – Set up

85. Performing a Titration – Set up Acid

86. Performing a Titration – Starting Amount

87. Performing a Titration – Set up Base

88. Performing a Titration - Titrating

89. Performing a Titration – End Point

90. How to solve a titration problem:
Start with the balanced equation for the neutralization reaction, and determine the chemically equivalent amounts of the acid and base.
Determine the moles of acid (or base) from the known solution used during the titration.
Determine the moles of solute of the unknown solution used during the titration.
Determine the molarity of the unknown solution.

91. Molarity and Titration
Determine the molarity of an acidic solution, 10 mL HCl, by titration
Titrate acid with a standard base solution
20.00 mL of 5.0 × 10−3 M NaOH was titrated
Write the balanced neutralization reaction equation.
HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)
1 mol mol mol mol

92. Molarity and Titration
Calculate the number of moles of NaOH used in the titration.
20.0 mL of 5.0 × 10−3 M NaOH is needed to reach the end point
mol of HCl = mol NaOH = 1.0 × 10−4 mol
Calculate the molarity of the HCl solution

93. Example Problem
In a titration, 27.4 mL of M Ba(OH)2 is added to a 20.0 mL sample of HCl solution of unknown concentration until the equivalence point is reached. What is the molarity of the acid solution?

94. Example Problem Solution
Given: mL of M Ba(OH)2
Unknown: ? M HCl of 20.0 mL
Solution:
Write balanced equation:
Ba(OH)2 + 2HCl BaCl2 + 2H2O
1 mol mol 1 mol 2 mol

95. 1. Calculate Moles of Given

96. 2. Write a mole ratio: moles of base used to moles of acid produced

97. 3. Calculate Unknown Molarity