1. Section 19.1 Acid – Base Theories
Chapter 19 Acids & Bases
Section Acid – Base Theories

2. Acids Bases vinegar  citrus fruits carbonated drinks  car battery
lemon juice  tea
Bases
calcium hydroxide in mortar  antacids
household cleaning agents

3. Properties of Acids Give foods a tart or sour taste
e.g. lemon & vinegar for example
Aqueous solutions of acids are electrolytes (conduct electricity)
Acids cause certain chemical indicators to change color.
Acid + Base Salt + water

4. Properties of Bases Bases have a bitter taste e.g. soap
Bases have a slippery feel
Aqueous solutions of bases are electrolytes (conduct electricity)
Bases cause certain chemical indicators to change color.
Acid + Base Salt + water

5. Arrhenius Acids & Bases
In 1887, Swedish chemist Svante Arrhenius
proposed a revolutionary way of defining and thinking about acids and bases
Acids are hydrogen-containing compounds that ionize to yield hydrogen ions (H+) in aqueous solution.
Bases are compounds that ionize to yield hydroxide ions (OH-) in aqueous solution

6. Monoprotic acids: acids that contain one
ionizable hydrogen
HNO3 – nitric acid
Diprotic acids: acids that contain two ionizable
hydrogens
H2SO4 – sulfuric acid
Triprotic acids: acids that contain three
ionizable hydrogens
H3PO4 – phosphoric acid

7. Not all compounds that contain hydrogen are acids
e.g. CH4 – methane has weak polar C – H bonds and no ionizable hydrogens. Not an acid.
Not all hydrogens in an acid may be released as hydrogen ions.
Only hydrogens in very polar bonds are ionizable. In the case where hydrogen is joined to a very electronegative element.
e.g. HCl hydrogen chloride very polar covalent molecule

8. When HCl dissolves in water, it releases hydrogen ions because the hydrogen ions are stabilized by solvation.
H2O
H – Cl (g) H+ (aq) + Cl- (aq)
Hydrogen Hydrogen Chloride
chloride ion ion
Ionizes to form an aqueous solution of hydronium ions and chloride ions
HCl H2O H3O Cl-

9. Ethanoic acid CH3COOH is a monoprotic acid due to its structure
H O
H C C O H H
The three H attached to the carbon are in weak polar bonds. They do not ionize.
Only the H bonded to the highly electronegative O can be ionized

10. NaOH (s) Na+ (aq) + OH- (aq)
Sodium hydroxide dissociates into sodium ions and hydroxide ions in aqueous solution.
H2O
NaOH (s) Na+ (aq) OH- (aq)
Sodium Sodium Hydroxide
Hydroxide Ion ion
Potassium hydroxide dissociates into potassium ions and hydroxide ions in aqueous solution.
H2O
KOH (s) K+ (aq) OH- (aq)
Potassium Potassium Hydroxide
Hydroxide Ion ion

11. Arrhenius Bases
Group one, the alkali metals, react with water to produce solutions that are basic.
Group one metals are very soluble in water and can produce concentrated solutions.
Group two metals are not very soluble in water. Their solutions are always very dilute.

12. Bronsted – Lowry Acids and Bases
The Bronsted – Lowry theory defines
Acid: a hydrogen-ion donor
Base: a hydrogen-ion acceptor
All acids and bases included in the Arrhenius theory are also acids and bases according to the Bronsted-Lowry theory.

13. Ammonia is the hydrogen-ion acceptor therefore it is a base.
NH3 (aq) H2O (l) NH4+ (aq) OH- (aq)
base acid conjugate conjugate
acid base
Ammonia is the hydrogen-ion acceptor
therefore it is a base.
Water is the hydrogen-ion donor and
therefore it is an acid.
Hydrogen ions are transferred from water to ammonia, which causes the hydroxide-ion concentration to be greater than it is in pure water.
When ammonia dissolves and reacts with water, NH4+ is the conjugate acid of the base NH3.
OH- is the conjugate base of acid H2O

14. Conjugate Acids and Bases
HCl (g) + H2O (l) H3O+ (aq) Cl - (aq)
acid base conjugate conjugate
acid base
HCl is the hydrogen-ion donor – thus it is an acid.
Water is the hydrogen-ion acceptor – thus it is a base

15. Conjugate Acid – Base Pair
Conjugate acid: the particle formed when a base gains a hydrogen ion
Conjugate base: the particle that remains when an acid has donated a hydrogen ion.
Conjugate acids and bases are always paired with a base or an acid, respectively.
Conjugate acid-base pairs consists of two substances related by the loss or gain of a single hydrogen ion.

16. Common Conjugate Acid – Base Pairs
Conjugate Base
HCl
Cl-
H2SO4
HSO4-
H3O+
H2O
SO42-
CH3COOH
CH3COO-
H2CO3
HCO3-
CO32-
NH4+
NH3
OH-

17. Bronsted – Lowry Acids and Bases
A water molecule that gains a hydrogen ion becomes a positively charged hydronium ion (H3O+)
Amphoteric – a substance that can act as both an acid and a base
e.g. water
H2SO4 + H2O H3O HSO4-
NH H2O NH OH-

18. Lewis Acids and Bases Gilbert Lewis proposed a third Acid Base theory
Acid – accepts a pair of electrons during a reaction
Base – donates a pair of electrons during a reaction
Concept is more general than either the Arrhenius theory or the Bronsted-Lowry theory.

19. Lewis Acids and Bases
Lewis Acid – a substance that can accept a pair of electrons to form a covalent bond.
Lewis Base – a substance that can donate a pair of electrons to form a covalent bond. ..
H : O – H :O: H H
Lewis Lewis
Acid Base

20. Electron-pair acceptor
Acid Base Definitions
Type
Acid
Base
Arrhenius
H+ producer
OH- producer
Bronsted - Lowry
H+ donor
H+ acceptor
Lewis
Electron-pair acceptor
Electron-pair donor

21. End of Section 19.1

22. Section 19.2 Hydrogen Ions & Acidity

23. Hydrogen Ions From Water
A water molecule that loses a hydrogen ion becomes a negatively charged hydroxide ion
A water molecule that gains a hydrogen ion becomes a positively charged hydronium ion
H2O (l) OH- (aq) H+ (aq)
Hydroxide ion Hydrogen ion
Self ionization of water – the reaction in which water molecules produce ions

24. Self Ionization of Water
Hydrogen ions in aqueous solution have several names.
Some chemists call them protons
Some chemists call them hydrogen ions or hydronium ions.
For our purposes, either H+ or H3O+ will represent hydrogen ions in aqueous solution.
H2O + H2O H3O OH-

25. In pure water at 25˚C, the equilibrium concentration of hydrogen ions and hydroxide ions are each only 1 x 10-7.
In other words the concentration of OH- and H+ are equal in pure water
Any aqueous solution in which H+ and OH- are equal is a neutral solution.
When [H+] increases [OH-] decreases
When [H+] decreases [OH-] increases

26. For aqueous solutions, the product of the hydrogen ion concentration and the hydroxide ion concentration equals 1.0 x 10-14
[H+] x [OH-] = 1 x 10-14
1 x x 1 x 10-7 = 1 x 10-14
This equation is true for all dilute aqueous solutions at 25˚C.
Ion-Product Constant for Water (Kw)
the product of the concentrations of the hydrogen ions and hydroxide ions in water
Kw = [H+] x [OH-] = 1.0 x 10-14

27. But not all solutions are neutral
When some substances dissolve in water, they release hydrogen ions.
When hydrogen chloride dissolves in water, it forms hydrochloric acid.
H2O
HCl (g) H+ (aq) + Cl- (aq)

28. Ion Product Constant for Water
In the previous HCl solution, the hydrogen-ion concentration is greater than the hydroxide-ion concentration.
Acidic Solution:
A solution in which [H+] is greater than [OH-].
The [H+] of an acidic solution is greater than 1 x 10-7 M

29. NaOH(s) Na+(aq) + OH-(aq)
When sodium hydroxide dissolves in water, it forms hydroxide ions in solution.
H20
NaOH(s) Na+(aq) + OH-(aq)
In the above solution, the hydrogen-ion concentration is less than the hydroxide-ion concentration.
Basic Solution:
A solution in which [H+] is less than [OH-]
The [H+] of a basic solution is less than 1 x 10-7
Basic solutions are also known as alkaline solutions.

30. The pH Concept
The pH scale was proposed by Danish Scientist Soren Sorensen in 1909.
The pH scale is used to express [H+]
Strongly Neutral Strongly
Acidic Basic

31. Calculating pH pH = - log [H+]
The pH of a solution is the negative logarithm of the hydrogen-ion concentration.
pH = - log [H+]

32. In neutral solution, the [H+] = 1 x 10-7M
pH = - log [H+]
pH = - log (1 x 10-7)
pH = 7

33. Classifying Solutions
A solution in which [H+] is greater than 1 x 10-7 has a pH less than 7 is called acidic.
A solution in which [H+] is less than 1 x 10-7 has a pH greater than 7 is called basic.
The pH of pure water or a neutral aqueous solution is 7
Acidic solution: pH < 7.0 [H+] > 1 x 10-7 M
Neutral solution: pH = 7.0 [H+] = 1 x 10-7 M
Basic solution: pH > 7.0 [H+] < 1 x 10-7 M

35. pH can be read from the value of [H+] if it is written in scientific notation and has a coefficient of 1.
Then the pH of the solution equals the exponent, with the sign changed from minus to plus
e.g. [H+] = 1 x has a pH of 2
e.g. [H+] = 1 x has a pH of 13

36. If the pH is an integer, it is also possible to directly write the value of [H+].
pH = 9 ; then [H+] = 1 x 10-9 M
pH = 4 ; then [H+] = 1 x 10-4M

37. A neutral solution has a pOH of 7
The pOH of a solution equals the negative logarithm of the hydroxide-ion concentration
pOH = - log [OH-]
A neutral solution has a pOH of 7
Acidic solution: pOH > 7 [OH-] < 1 x 10-7 M
Neutral solution: pOH = 7 [OH-] = 1 x 10-7 M
Basic solution: pOH < 7 [OH-] > 1 x 10-7 M

38. pH and pOH Relationship
pOH + pH = 14
pH = 14 – pOH
pOH = 14 – pH

40. Problem Example
Colas are slightly acidic. If the [H+] in a solution is 1.0 X M , is the solution acidic, basic or neutral. What is the [OH-] of this solution?
[H+] = 1.0 X M which is greater than 1.0 X M
so the solution is acidic
Kw = [OH-] x [H+] = 1.0 X
[OH-] = 1.0 X ÷ [H+]
[OH-] = 1.0 X ÷ 1.0 X
[OH-] = 1.0 X M

41. Problem Example
What is the pH of a solution with a hydrogen-ion concentration of 4.2 x M?
pH = - log [H+]
pH = - log (4.2 x )
pH = 9.38

42. Using calculator find the anti log of - 6.35
Problem Example
pH of an unknown solution is What is its hydrogen-ion concentration?
pH = -log [H+]
6.35 = -log [H+]
= log [H+]
Using calculator find the anti log of
[H+] = 4.5 x M

43. Problem Example
What is the pH of a solution if the [OH-] = 4.0 X10 – 11 M?
Kw = [H+] x [OH-] = 1 x
[H+] = 1 x ÷ [OH-]
[H+] = 1 x ÷ 4.0 x
[H+] = 2.5 x M

44. Problem Example (con’t)
What is the pH of a solution if [OH-] = 4.0 X M?
pH = - log [H+]
pH = - log (2.5 x )
pH = 3.60

45. Acid – Base Indicators
Indicator - is an acid or a base that undergoes dissociation in a known pH range
An indicator is a valuable tool for measuring pH because its acid form and base form have different color in solution.
For each indicator, the change from dominating acid form to dominating base form occurs in a narrow range of approximately two pH units.
Within this range, the color of the solution is a mixture of the colors of the acid and the base forms.
Knowing the pH range over which this color change occurs, can give you a rough estimate of the pH of the solution.

47. End of Section 19.2

48. Section 19.3 Hydrogen Ions & Acidity

49. Strong Acids
Acids are classified as strong or weak depending on the degree to which they ionize in water.
In general, strong acids are completely ionized in aqueous solution.
HNO nitric acid HCl hydrochloric acid H2SO sulfuric acid HClO perchloric acid HBr hydrobromic acid HI hydroiodic acid
HCl(g) + H2O(l) H3O +(aq) + Cl ¯(aq)

50. Weak Acids Weak acids ionize only slightly in aqueous solution.
Some Weak Acids
Acetic Acid CH3COOH
Boric Acid H3BO3 (all three are weak)
Phosphoric Acid H3PO4 (all three are weak)
Sulfuric Acid HSO4- (first ionization is strong)
CH3COOH(aq) + H2O(l) H3O+(aq) + CH3COO ¯ (aq)
ethanoic acid water hydronium ethanoate
ion ion

51. Acid Strength
A strong acid completely dissociates in water ([H3O+] is high).
A weak acid remains largely undissociated.
([H3O+] is low).

52. Equilibrium Constant (Keq)
Write the equilibrium-constant expression from the balanced chemical equation.
CH3COOH(aq) + H2O(l) H3O+(aq) + CH3COO ¯(aq)
Keq = [H3O+] x [ CH3COO ¯]
[H3COOH] x [H2O] [H2O] constant in dilute solutions

53. Acid Dissociation Constant (Ka)
Ka = Ratio of the concentration of the dissociated
form of an acid to the concentration of the
undissociated form.
CH3COOH (aq) + H2O (l) H3O +(aq) + CH3COO ¯(aq)
Acid Dissociation Constant
Ka = [H3O+] x [ CH3COO¯]
[CH3COOH]

54. Acid Dissociation Constant (Ka)
Acid dissociation constant reflects the fraction of an acid in the ionized form.
(Ka sometimes called ionization constant)
If the value of the Ka is small, then the degree of dissociation or ionization of the acid in the solution is small.
Weak acids – small Ka values
Stronger the acid – larger the Ka

55. Acid Dissociation Constant (Ka)
Nitrous acid (HNO2) has a Ka of 4.4 x 10¯4
Acetic acid (CH3COOH) has a Ka of 1.8 x 10¯5
Nitrous acid is more ionized in solution and a stronger
acid

56. Acids Strong Acids Have high [H3O+] Large dissociation constant (Ka)
Weak Acids
Have low [H3O+]
Small dissociation constant (Ka)

57. Acids Diprotic and triprotic acids lose their hydrogens one at a time.
Each ionization reaction has a separate dissociation constant.
H3PO4 – 3 separate dissociation constants.

58. A 0. 012 M solution of formic acid HCOOH is partially ionized
A M solution of formic acid HCOOH is partially ionized. [H+] is 1.02 x 10-4 M. what is the acid dissociation constant (Ka) of formic acid?
Concentration
[HCOOH]
[H+]
[HCOO¯ ]
Initial
0.012
Change
– 1.02 x
1.02 x
1.02 x 10 – 4
Equilibrium
Substitute the equilibrium values into the expression for Ka :
[H+]
[HCOO¯ ]
(1.02 x 10 – 4)
(1.02 x 10 – 4)
Ka = = =
8.68 x10 - 7
[HCOOH]
Classwork page: # 22 & 23

59. Strong Bases
Bases are classified as strong or weak depending on the degree to which they ionize in water.
In general, strong bases are completely ionized in aqueous solution.
Ca(OH) calcium hydroxide NaOH sodium hydroxide KOH potassium hydroxide
NaOH(s) + H2O(l) Na +(aq) + OH ¯(aq)

60. Weak bases ionize only slightly in aqueous solution.
Some Weak bases
Ammonia NH3
Methylamine CH3NH2
NH3 (aq) + H2O(l) NH4+(aq) OH ¯ (aq) ammonia water ammonium ion hydroxide ion

61. Base Dissociation Constant (Kb)
Strong bases dissociate completely into metal ions and hydroxide ions in aqueous solution.
Some strong bases are not very soluble in water (calcium hydroxide and magnesium hydroxide)
Small amounts that do not dissolve dissociate completely
Weak bases react with water to form the hydroxide ion and the conjugate acid of the base.
NH3(aq) + H2O(l) NH4+ (aq) + OH¯ (aq)
Ammonia Water Ammonium Ion Hydroxide ion

62. Base Dissociation Constant (Kb)
NH3(aq) + H2O(l) NH4+(aq) + OH ¯(aq)
Ammonia Water Ammonium Ion Hydroxide ion
Only about 1% of ammonia is present as NH4+
Base Dissociation Constant (Kb)
Kb = [NH4+] x [OH¯ ]
[NH3]

63. Concentration and Strength
The words concentrated and dilute indicate how much of an acid or base is dissolved in solution.
Number of moles of the acid or base in a given volume.
The words strong and weak refer to the extent of ionization or dissociation of an acid or base.
How many of the particles ionize or dissociate into ions
A sample of HCl added to a large volume of water becomes more dilute, but it is still a strong acid.
Vinegar is a dilute solution of a weak acid.

64. End of section 19.3

65. Section 19.4 Neutralization Reactions

66. Acid – Base Reactions
If you mix a solution of a strong acid containing hydronium ions with a solution of a strong base that has an equal number of hydroxide ions, a neutral solution results.
Final solution has properties that are characteristic of neither an acidic nor a basic solution.
HCl (aq) + NaOH (aq) NaCl (aq) + H2O (l)
H2SO4 (aq) + 2KOH (aq) K2SO4 (aq) + H2O (l)

67. Neutralization Reactions
Reactions of weak acids and weak bases do not
usually produce a neutral solution.
In general, reactions with which an acid and a base
react in an aqueous solution to produce a salt and
water are called neutralization reactions.

68. Making Salts Prepare potassium chloride by mixing equal molar
quantities of hydrochloric acid and potassium
hydroxide.
HCl + KOH KCl + H2O
Heating the solution to evaporate the water will leave
the salt potassium chloride.
In general, the reaction of an acid with a base
produced water and salt

69. Titration The number of moles of hydrogen ions provided by
the acid are equivalent to the number of hydroxide
ions provided by the base.
HCl (aq) + NaOH (aq) NaCl (aq) + H2O (l)
1 mole mole mole mole
H2SO4 (aq) + 2NaOH (aq) Na2SO4 (aq) + 2H2O (l)
1 mole mole mole mole
When an acid & base are mixed, the Equivalence
point is when the number of moles of hydrogen ions
equals the number of moles of hydroxide ions.

70. Sample Problem How many moles of sulfuric acid are required to
neutralize 0.50 mol of sodium hydroxide?
H2SO4 (aq) + 2NaOH (aq) Na2SO4 (aq) + 2H2O (l)

71. Practice Problem How many moles of potassium hydroxide are needed
to completely neutralize 1.56 mol of phosphoric acid?
H3PO4 (aq) + 3KOH (aq) K3PO4 (aq) + 3H2O (l)

72. Titration You can determine the concentration of acid or base
in a solution by performing a neutralization reaction.
You must use an appropriate acid-base indicator to
show when neutralization has occurred.
In the lab, typically phenolphthalein for acid base
neutralization reactions.
Solutions that contain phenolphthalein turn from
colorless to deep pink as the pH of the solution
changes from acidic to basic.

73. Titration Measured volume of an acid solution of unknown
concentration is added to a flask

74. Titration Several drops of the indicator are added to the
solution while the flask is swirled

75. Titration Measured volumes of the base of known concentration
are mixed into the acid until the
indicator just barely changes color.

76. Titration Titration – the process of adding a known amount of
solution of known concentration to determine the
concentration of another solution.
Standard solution – the solution of known
concentration
End point – the point at which the indicator changes
color, the point of neutralization
You can also use titration to find the concentration of
a base using a standard acid.
Equivalence point – the point in a titration where the
number of moles of hydrogen ions = number of
moles of hydroxide ions.

77. A 25ml solution of H2SO4 is completely neutralized
by 18ml of 1.0M NaOH. What is the concentration of
the H2SO4 solution?
H2SO4 (aq) + 2NaOH (aq) Na2SO4 (aq) + 2H2O (l)

78. Practice Problem How many milliliters of 0.45M HCl will neutralize
25.0ml of 1.00M KOH?

79. Practice Problem
What is the molarity of H3PO4 if 15.0 ml is completely neutralized by 38.5 ml of M Ca(OH)2?

80. End of section 19.4

81. Section 19.5 Salts in Solution

82. Salt Hydrolysis A salt consists of an anion from an acid and a cation
from a base.
The salt forms as a result of a neutralization reaction
Although solutions of many salts are neutral, some
are acidic and others are basic.

83. Salt Hydrolysis Salt Hydrolysis – the cations or anions of
a dissociated salt remove hydrogen ions from or
donate hydrogen ion to water.
Hydrolyzing salts are usually derived from a strong
acid and weak base or from a weak acid and a
strong base.
In general, salts that produce acidic solutions contain
positive ions that release protons to water.
Salts that produce basic solutions contain negative
ions that attract protons from water.

84. Salt Hydrolysis CH3COONa (aq) CH3COO ¯ (aq) + Na+ (aq)
Sodium ethanoate ethanoate ion sodium ion
CH3COONa is the salt from a weak acid CH3COOH
and a strong base NaOH
In solution, the salt is completely ionized.

85. Salt Hydrolysis
Salt Hydrolysis – the cations or anions of a dissociated salt remove hydrogen ions from or donate hydrogen ion to water.
CH3COO¯ (aq) + H2O (l) CH3COOH (aq) + OH ¯(aq)
BL base BL acid makes
hydrogen-ion hydrogen-ion solution
acceptor donor basic
This process is called hydrolysis because it splits a
hydrogen ion off a water molecule.
Resulting solution contains a hydroxide-ion
concentration greater than the hydrogen-ion
concentration. Thus the solution is basic

86. Salt Hydrolysis NH4Cl (aq) NH4+ (aq) + Cl ¯ (aq)
Ammonium Ammonium ion Chloride ion
chloride
NH4Cl is the salt from a strong acid (hydrochloric acid, HCl) and a weak base (ammonia, NH3)
In solution the salt is completely ionized.

87. Salt Hydrolysis NH4+ (aq) + H2O (l) NH3 (aq) + H3O+ (aq)
BL acid BL base
Hydrogen-ion Hydrogen-ion
donor acceptor
This process is also called hydrolysis because it splits a hydrogen ion off a water molecule.
Resulting solution contains a hydrogen-ion concentration greater than the hydroxide-ion concentration. Thus the solution is acidic

88. Salt Hydrolysis Equivalence Point Strong Acid Strong Base pH= 7
neutral
Weak Acid
Strong Base
pH > 7
basic
Strong Acid
Weak Base
pH < 7
acidic
Equivalence point – the point in a titration where the number of moles of hydrogen ions = number of moles of hydroxide ions

89. Buffers
Buffer – a solution in which the pH remains relatively constant when small amounts of acid or base are added.
A buffer is a solution of a weak acid and one of its salts, or a solution of a weak base and one of its salts.
A buffer solution is better able to resist drastic changes in pH than is pure water.

90. Buffers
A solution of ethanoic acid (CH3COOH) and sodium ethanoate (CH3COONa) is an example of a typical buffer.
CH3COOH and CH3COO - (source is the completely ionized CH3COONa) act as reservoirs of neutralizing power.

91. Buffers CH3COO - (aq) + H+ (aq) CH3COOH (aq)
ethanoate ion Hydrogen ion ethanoic acid
When an acid is added to the solution, the ethanoate ions act as a hydrogen-ion sponge.
CH3COOH (aq) + OH - (aq) CH3COO - (aq) + H2O (l)
Ethanoic acid hydroxide ion ethanoate ion water
When a base is added to the solution, the ethanoic acid and the hydroxide ions react to produce water and the ethanoate ion.

92. Buffers
The ethanoate ion is not strong enough base to accept hydrogen ions from water extensively.
The buffer solution cannot control the pH when too much acid is added, because no more ethanoate ions are present to accept hydrogen ions.
Buffer also become ineffective when too much base is added. No more ethanoic acid molecules are present to donate hydrogen ions.

93. Buffers
When too much acid or base is added, the buffer capacity is exceeded.
Buffer capacity – the amount of acid or base that can be added to a buffer solution before a significant change in pH occurs.

94. Buffers
When a base is added to a buffered solution, the acidic form removes hydroxide ions from the solution.
When an acid is added to a buffered solution, the basic form removes hydrogen ions from the solution.

95. Buffers & Your Blood
Your body function properly only when the pH of your blood lies between 7.35 and 7.45
Your blood contains buffers (hydrogen carbonate ions and carbonic acid)
HCO3- (aq) H+ (aq) H2CO3 (aq)
Hydrogen Hydrogen ion Carbonic acid
carbonate ion
As long as there are hydrogen carbonate ions available, the excess hydrogen ions are removed, and the pH of the blood changes very little.

96. End of Section Chapter 19