1. Chapter 19 Acids, Bases, and Salts

2. Acids and Bases 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
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 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
Chemists recognized the properties of acids and bases, but were unable to propose a theory to explain their behavior.
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. Arrhenius Acids
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. Arrhenius Acids Not all compounds that contain hydrogen are acids
Ex. 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.
Ex. HCl hydrogen chloride very polar covalent molecule

8. Arrhenius Acids
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. Arrhenius Acids
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. Arrhenius Bases
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 sodium ions and hydroxide ions in aqueous solution.
H2O
KOH (s) K+ (aq) + OH- (aq)
Potassium Potassium Hydroxide
Hydroxide Ion ion

11. Arrhenius Bases
Group IA, the alkali metals, react with water to produce solutions that are basic.
Group IA metals are very soluble in water and can produce concentrated solutions.
Group 2A metals are not very soluble in water. Their solutions are always very dilute.

12. Bronsted-Lowry Acids and Bases
Arrhenius’ definition of acids and bases is not a very comprehensive one.
If defines acids and bases narrowly and does not include certain substances that have acidic or basic properties.
Na2CO3 (aq) is basic

13. Bronsted-Lowry Acids and Bases
The Bronste-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.

14. Ammonia as a Base Bronsted-Lowry Theory
NH3 (aq) + H2O (l) NH4+ (aq) + OH- (aq)
ammonia is the hydrogen-ion acceptor and therefore a BL base
water is the hydrogen-ion donor and therefore a BL acid.
Hydrogen ions are transferred from water to ammonia, which causes the hydroxide-ion concentration to be greater than it is in pure water.

15. Conjugate Acids and Bases
NH3 (aq) + H2O (l) NH4+ (aq) + OH- (aq)
base acid conjugate acid conjugate base
When ammonia dissolves and reacts with water, NH4+ is the conjugate acid of the base NH3.
OH- is the conjugate base of acid H2O

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

17. 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.

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

19. 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
Ex: water
H2SO4 + H2O H3O HSO4-
NH H2O NH OH-

20. 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.

21. 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

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

23. End of Section 19.1

24. Hydrogen Ions From Water
Water molecules are highly polar and are in continuous motion.
Occasionally, the collisions between water molecules are energetic enough to transfer a hydrogen ion from one water molecule to another.
Self ionization of water – the reaction in which water molecules produce ions

25. 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 Hydroxide ion
Self ionization of water – the reaction in which water molecules produce ions

26. 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-

27. Self Ionization of Water
The self-ionization of water occurs to a very small extent.
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

28. Neutral Solutions
Any aqueous solution in which H+ and OH- are equal is a neutral solution.

29. Ion Product Constant for Water
When [H+] increases [OH-] decreases
When [H+] decreases [OH-] increases
LeChatelier’s principle – when a stress is applied to a system in dynamic equilibrium, the system changes in a way that relieves the stress
If additional ions (either H+ or OH-) are added to a solution, the equilibrium shifts. The concentration of the other type of ion decreases. More water molecules are formed in the process.
H+ (aq) + OH- (aq) H2O (l)

30. Ion Product Constant for Water
For aqueous solutions, the product of the hydrogen ion concentration and the hydroxide ion concentration equals 1.0 x 10-14
[H+] x [OH-] = 1.0 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

31. Ion Product Constant for Water
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)

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

33. Ion Product Constant for Water
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 – one 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.

34. 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

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

36. Calculating pH In neutral solution, the [H+] = 1 x 10-7M. The pH is 7
pH = -log[H+]
pH = -log(1 x 10-7)
pH = -(log 1 + log 10-7)
pH = -( )
pH = 7.0

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

39. Calculating pH
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
[H+] = 1 x 10-2 has a pH of 2.0
[H+] = 1 x has a pH of 13.0

40. Calculating pH
If the pH is an integer, it is also possible to directly write the value of [H+].
pH = 9.0 then [H+] of 1 x 10-9M
pH = 4 then [H+] = 1 x 10-4M

41. A neutral solution has a pOH of 7
Calculating pOH
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 > 7.0 [OH-] < 1 x 10-7M
Neutral solution: pOH = 7.0 [OH-] equals 1 x 10-7M
Basic solution: pOH < 7.0 [OH-] > 1 x 10-7

42. pH and pOH Relationship
pOH + pH – 14
pH= 14 – pOH
pOH = 14 - pH

44. pH Significant Figures
For pH calculation, you should express the hydrogen-ion concentration in scientific notation
[H+] = M should be written 1.0 x 10-3
0.0010M has two sig figs
Write pH = 3.00 with 2 zeros to the right of the decimal place representing the 2 sig figs

45. Problem Example
Colas are slightly acidic. If the [H+] in a solution is 1.0 X 10-5 M , is the solution acidic, basic or neutral. What is the [OH-] of this solution?
[H+] = 1.0 X 10-5 M which is greater than 1.0 X 10-7 M so solution is acidic
Kw = [OH-] x [H+] = 1.0 X 10-14
[OH-] = 1.0 X / [H+]
[OH-] = 1.0 X / 1.0 X 10-5
[OH-] = 1.0 X 10-9

46. 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 10-10)
pH = -(9.3765)
pH = 9.38

47. Using calculator find the antilog of -6.35
Problem Example
pH of an unknown solution is What is its hydrogen-ion concentration?
pH = -log [H+]
6.35 = -log [H+]
-6.35 = log [H+]
Using calculator find the antilog of -6.35
4.5 x 10-7 M = [H+]

48. Problem Example What is the pH of a solution if the [OH-] =4.0X10-11M?
Kw = [H+] x [OH-]= 1 x 10-14
[H+] = 1 x / [OH-]
[H+] = 1 x / 4.0 x 10-11
[H+] =0.25 x 10-3 M
[H+] = 2.5 x 10-4 M

49. Problem Example (con’t)
What is the pH of a solution if the
[OH-] =4.0X10-11M?
pH = -log [H+]
pH = -log (2.5 x 10-4)
pH = - ( )
pH = 3.60

50. HIn (aq) H+ (aq) + In- (aq)
Acid-Base Indicators
Indicator - (HIn) is an acid or a base tht undergoes dissociation in a know pH range
An indicator is a valuable tool for measuring pH because its acid form and base form have different color in solution.
OH-
HIn (aq) H+ (aq) In- (aq)
acid H base
form form
The acid form dominates the dissociation equilibrium at low pH (high [H+]), and the base form dominates the equilibrium at high pH (high [OH-])

51. Acid-Base Indicators
For each indicator, the change from dominating acid from 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.

52. Indicator characteristics that limit their usefulness.
Acid-Base Indicators
Many different indicators are needed to span the entire pH spectrum.
Indicator characteristics that limit their usefulness.
Listed pH values of indicators are usually given for 25ºC. At other temperatures, an indicator may change color at a different pH.
If the solution being tested is not colorless, the color of the indicator may be distorted.
Dissolved salts in a solution may also affect the indicator’s dissociation.
Using indicator strips can help overcome these problems.

54. A pH meter makes rapid, accurate pH measurements.
pH Meters
A pH meter makes rapid, accurate pH measurements.
often easier to use than liquid indicators or indicator strips.
Measurements of pH obtained with a pH meter are typically accurate to within 0.01 pH unit of the true pH.
Color and cloudiness of the unknown solution do not affect the accuracy of the pH value
If the solution being tested is not colorless, the color of the indicator may be distorted.

55. End of section 19.2

56. HCl(g) + H2O(l) H3O+(aq) + Cl-(aq)
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.
HNO3 - nitric acid HCl - hydrochloric acid H2SO4 - sulfuric acid HClO4 - perchloric acid HBr - hydrobromic acid HI - hydroiodic acid
HCl(g) + H2O(l) H3O+(aq) + Cl-(aq)

57. Weak Acids Weak acids ionize only slightly in aqueous solution.
Some Weak Acids
Acetic Acid H3COOH
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

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

59. 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

60. Acid Dissociation Constant (Ka)
Ka = Ratio of the concentration of the dissociated form of an acid to the concentration of the undissociated form.
H3COOH(aq) + H2O(l) H3O+(aq) + CH3COO-(aq)
Acid Dissociation Constant
Ka = [H3O+] x [ CH3COO- ]
[CH3COOH]

61. 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

62. 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

63. Acids Strong Acids Have high [H3O+] Large dissociation constant
Weak Acids
Have low [H3O+]
Small dissociation constant

64. 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.

65. 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

66. 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+
Equilibrium Constant
Keq = [NH4+] x [OH- ]
[NH3] x [H2O] [H2O] constant in dilute solutions
Base Dissociation Constant
Kb = [NH4+] x [OH- ]
[NH3]

67. 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.

68. End of section 19.3

69. 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)

70. 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.

71. Making Salts
Prepare potassium chloride by mixing equal molar quantities of hydrochloric acid and potassium hydroxide.
HCl + KOH KCl + H20
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

72. 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) + H20 (l)
1 mole mole mole mole
H2SO4(aq) + 2NaOH(aq) Na2SO4(aq) + 2H20 (l)
1 mole mole mole mole
When and acid & base are mixed, the Equivalence point is when the number of moles of hydrogen ions equals the number of moles of hydroxide ions.

73. H2SO4(aq) + 2NaOH(aq) Na2SO4(aq) + 2H20 (l)
Sample Problem
How many moles of sulfuric acid are required to neutralize 0.50 mol of sodium hydroxide?
H2SO4(aq) + 2NaOH(aq) Na2SO4(aq) + 2H20 (l)
Mole ratio of H2SO4 to NaOH is 1:2
0.50 mol NaOH mol H2SO = mol H2SO4
2 mol NaOH

74. H3PO4(aq) + 3KOH(aq) K3PO4(aq) + 3H2O(l)
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)
1.56 mol H3PO mol KOH = mol H3PO4
1 mol H3PO4

75. 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.

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

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

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

79. 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
You can also use titration to find the concentration of a base using a standard acid.

80. 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
Equivalence point – the point in a titration where the number of moles of hydrogen ions = number of moles of hydroxide ions..

81. H2SO4(aq) + 2NaOH(aq) Na2SO4(aq) + 2H20 (l)
Sample Problem
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) + 2H20 (l)
0.018 L NaOH 1.0 mol NaOH 1 mol H2SO =
1L NaOH mol NaOH L
0.36M H2SO4

82. Practice Problem
How many milliliters of 0.45M HCl will neutralize 25.0ml of 1.00M KOH?
HCl + KOH H2O + KCl
0.025 L KOH 1.0 mol KOH mol HCl
1L KOH mol KOH
1 L HCl ml HCl = 56 ml HCl
0.45 mol HCl L HCl

83. Practice Problem
What is the molarity of H3PO4 if 15.0 ml is completely neutralized by 38.5 ml of M?
H3PO4 + 3NaOH H2O + Na3PO4
L NaOH mol NaOH 1 mol H3PO4
1L NaOH mol NaOH
= M H3PO4
0.015L H3PO4

84. End of section 19.4

85. 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. ..

86. 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.

87. 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.

88. 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

89. 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.

90. NH4+(aq) + H2O(l) NH3(aq) + H3O+(aq)
Salt Hydrolysis
NH4+(aq) + H2O(l) NH3(aq) + H3O+(aq)
BL acid BL base makes
hydrogen-ion hydrogen-ion solution
donor acceptor acidic
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

91. 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

92. 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.

93. 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.

94. CH3COO-(aq) + H+(aq) CH3COOH (aq)
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.

95. 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.

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

97. 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.

98. HCO3- (aq) + H+ (aq) H2CO3 (aq)
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 ions 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.

99. End of Section Chapter 19