2. Acids and Bases
Acids
• vinegar • citrus fruits
• carbonated drinks • car battery
• lemon juice • tea
Bases
• calcium hydroxide in mortar • antacids
• household cleaning agents
3. Properties of Acids
1. Give foods a tart or sour taste
a)lemon & vinegar for example
2. Aqueous solutions of acids are electrolytes
(conduct electricity)
3. Acids cause certain chemical indicators to
change color.
4. Acid + Base Salt + water
4. Properties of Bases
1. Bases have a bitter taste
a)soap
2. Bases have a slippery feel
3. Aqueous solutions of bases are electrolytes
(conduct electricity)
4. Bases cause certain chemical indicators to
change color.
5. 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 + H 2O H 3O + + 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
Acid Base
HCl Cl-
H2SO4 HSO4-
H 3O + H 2O
HSO4- SO42-
CH3COOH CH3COO-
H2CO3 HCO3-
HCO3- CO32-
NH4+ NH3
H 2O 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-
NH3 + H2O NH4+ + 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. Acid Base Definitions
Type Acid Base
Arrhenius H+ producer OH- producer
Bronsted Lowry H+ H+ acceptor
Electron-pair Electron-pair
Lewis acceptor donor
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
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
+ - -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.
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+]
1 2 3 4 5 6 7 8 9 10 11 12 13 14
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 = -(0.0 + -7.0)
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
38.
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 10-13 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. 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
44. pH Significant Figures
For pH calculation, you should express the
hydrogen-ion concentration in scientific
notation
[H+] = 0.0010M 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 10-14 / [H+]
[OH-] = 1.0 X 10-14 / 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 10-10 M?
pH = -log [H+]
pH = -log (4.2 x 10-10)
pH = -(9.3765)
pH = 9.38
47. Problem Example
pH of an unknown solution is 6.35. 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 10-14 / [OH-]
[H+] = 1 x 10-14 / 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 = - (-3.60205)
pH = 3.60
50. 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. 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.
53.
54. 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.
56. 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.
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.
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 1 mole 1 mole 1 mole
H2SO4(aq) + 2NaOH(aq) Na2SO4(aq) + 2H20 (l)
1 mole 2 mole 1 mole 2 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. 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 1 mol H2SO4 = 0.25 mol H2SO4
2 mol NaOH
74. 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 H3PO4 3 mol KOH = 4.68 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.
78. Titration
c. 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. 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 H2SO4 =
1L NaOH 2 mol NaOH 0.025L
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 1 mol HCl
1L KOH 1 mol KOH
1 L HCl 1000 ml HCl = 56 ml HCl
0.45 mol HCl 1 L HCl
83. Practice Problem
What is the molarity of H3PO4 if 15.0 ml is completely
neutralized by 38.5 ml of 0.150 M?
H3PO4 + 3NaOH 3H2O + Na3PO4
0.0385 L NaOH 0.150 mol NaOH 1 mol H3PO4
1L NaOH 3 mol NaOH
= 0.129M H3PO4
0.015L H3PO4
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. 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
pH= 7 neutral
Strong Base
Weak Acid
pH > 7 basic
Strong Base
Strong Acid
pH < 7 acidic
Weak Base
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. 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. 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.
