1. 9.1 Introduction to Acids and Bases
The earliest definition was given by Arrhenius:
An acid contains a hydrogen atom and dissolves in water to form a hydrogen ion, H+.
HCl(g)
H+(aq) Cl−(aq)
acid
A base contains hydroxide and dissolves in water to form OH−.
NaOH(s)
Na+(aq) OH− (aq)
base

2. 9.1 Introduction to Acids and Bases
The Arrhenius definition correctly predicts the behavior of many acids and bases.
However, this definition is limited and sometimes inaccurate.
For example, H+ does not exist in water. Instead, it reacts with water to form the hydronium ion, H3O+.
H+(aq) H2O(l)
H3O+(aq)
hydrogen ion:
does not really exist
in solution
hydronium ion:
actually present in
aqueous solution

3. 9.1 Introduction to Acids and Bases
The Brønsted–Lowry definition is more widely used:
A Brønsted–Lowry acid is a proton (H+) donor.
A Brønsted–Lowry base is a proton (H+) acceptor.
HCl(g) H2O(l)
H3O+(aq) Cl−(aq)
HCl is a Brønsted–Lowry acid.
H2O is a Brønsted–Lowry base.

4. 9.1 Introduction to Acids and Bases A. Brønsted–Lowry Acids
A Brønsted–Lowry acid must contain a hydrogen atom.
Common Brønsted–Lowry acids (HA):
HCl
hydrochloric acid
H2SO4
sulfuric acid H
HBr
hydrobromic acid O
acidic H
atom H C C O H
HNO3
nitric acid H
acetic acid

5. 9.1 Introduction to Acids and Bases A. Brønsted–Lowry Acids
The names of common acids are derived from the anions formed when they dissolve in water.
Anions whose names end in –ide, add the prefix hydro and change the –ide ending to –ic acid.

6. 9.1 Introduction to Acids and Bases A. Brønsted–Lowry Acids
Sulfur based polyatomic acids = -ur…
Phosphorous based polyatomic acids = -or…
Polyatomic anions whose names end in –ate, change the –ate ending to –ic acid.
Polyatomic anions whose names end in –ite, change the –ite ending to –ous acid.

7. Name the following acids
HBr
H3PO4 (Phosphate)
H2CO3 (Carbonate)
CH3CO2H (Acetate)
HNO2 (Nitrite)

8. 9.1 Introduction to Acids and Bases A. Brønsted–Lowry Acids
A monoprotic acid contains one acidic proton.
A diprotic acid contains two acidic protons.
A triprotic acid contains three acidic protons.
A Brønsted–Lowry acid may be neutral or it may carry a net positive or negative charge.

9. Monoprotic, diprotic, or triprotic?
HBr
H3PO4
H2CO3
CH3CO2H
HNO2

10. Which of the following can be a Brønsted–Lowry Acid?HF
HSO3-
Cl2
SO32-
H2PO4-
Cl-

11. 9.1 Introduction to Acids and Bases B. Brønsted–Lowry Bases
A Brønsted–Lowry base is a proton acceptor, so it must be able to form a bond to a proton.
A base must contain a lone pair of electrons that can be used to form a new bond to the proton.

12. 9.1 Introduction to Acids and Bases B. Brønsted–Lowry Bases
Common Brønsted–Lowry Bases (B:)
Lone pairs make these
neutral compounds bases.
The OH− is the base in
each metal salt.

13. 9.2 The Reaction of a Brønsted–Lowry Acid with a Brønsted–Lowry Base+ B
A − + H B+

14. 9.2 The Reaction of a Brønsted–Lowry Acid with a Brønsted–Lowry Base+ B
A − + H B+
The product formed by loss of a proton from an acid is called its conjugate base.
The product formed by gain of a proton by a base is called its conjugate acid.

15. 9.2 The Reaction of a Brønsted–Lowry Acid with a Brønsted–Lowry Base+
H2O
Br− +
H3O+
acid
base
conjugate
base
conjugate
acid
The net charge must be the same on both sides of the equation.

16. Label the acid, the base, the conjugate acid and the conjugate base in each reaction
H2O HI  I H3O+
CH3COOH NH3  CH3COO NH4+
Br HNO3  HBr NO3-

17. 9.2 The Reaction of a Brønsted–Lowry Acid with a Brønsted–Lowry Base
When a species gains a proton (H+), it gains a +1 charge.
H2O
base
zero charge
H3O+
+1 charge
When a species loses a proton (H+), it effectively gains a −1 charge.
HBr
acid
zero charge
Br−
−1 charge

18. 9.2 The Reaction of a Brønsted–Lowry Acid with a Brønsted–Lowry Base
Amphoteric compound: A compound that contains
both a hydrogen atom and a lone pair of e−; it can
be either an acid or a base. + H
add H+ H O H H O H
H2O as a base
conjugate acid −
remove H+ H O H H O
H2O as an acid
conjugate base

19. 9.3 Acid and Base Strength A. Relating Acid and Base Strength
The splitting apart of a covalent molecule (or ionic compound) into its ions is called dissociation.
When a strong acid dissolves in water, 100% of the acid dissociates into ions.
HCl(g) H2O(l)
H3O+(aq) Cl−(aq)
A single reaction arrow is used, because the product is greatly favored at equilibrium.

20. 9.3 Acid and Base Strength A. Relating Acid and Base Strength
When a weak acid dissolves in water, only a small fraction of the acid dissociates into ions.
Unequal reaction arrows are used, because the reactants are usually favored at equilibrium.
CH3COOH(l) + H2O(l)
H3O+(aq) + CH3COO−(aq)

21. 9.3 Acid and Base Strength A. Relating Acid and Base Strength
A strong acid, HCl, is completely dissociated
into H3O+(aq) and Cl−(aq).
A weak acid, CH3COOH, contains mostly undissociated acid.

22. 9.3 Acid and Base Strength A. Relating Acid and Base Strength
When a strong base dissolves in water, 100% of the base dissociates into ions.
NaOH(s) H2O(l)
Na+(aq) −OH(aq)
Common strong bases are NaOH and KOH.
When a weak base dissolves in water, only a small fraction of the base dissociates into ions.
NH3(g) + H2O(l)
NH4+(aq) + −OH(aq)

23. 9.3 Acid and Base Strength A. Relating Acid and Base Strength
A strong base, NaOH, is completely dissociated
into Na+(aq) and −OH(aq).
A weak base contains
mostly undissociated
base, NH3.

24. 9.3 Acid and Base Strength A. Relating Acid and Base Strength
A strong acid readily donates a proton, forming
a weak conjugate base.
HCl
strong acid
Cl−
weak conjugate base
A strong base readily accepts a proton, forming
a weak conjugate acid.
OH−
strong base
H2O
weak conjugate acid

25. 9.3 Acid and Base Strength A. Relating Acid and Base Strength

26. Label the stronger acid in each pair
Label the stronger acid in each pair. Which acid has the stronger conjugate base?
H2SO4 OR H3PO4
HF OR HCl
H2CO3 OR NH4+
HCN OR HF

27. 9. 3. Acid and Base Strength B
9.3 Acid and Base Strength B. Using Acid Strength to Predict the Direction of Equilibrium
A Brønsted–Lowry acid–base reaction represents an equilibrium. H A + B
A − + H B+
acid
base
conjugate
base
conjugate
acid
The position of the equilibrium depends upon the strengths of the acids and bases.
The stronger acid reacts with the stronger base to form the weaker acid and the weaker base.

28. 9. 3. Acid and Base Strength B
9.3 Acid and Base Strength B. Using Acid Strength to Predict the Direction of Equilibrium
When the stronger acid and base are the reactants on the left side, the reaction readily occurs and the reaction proceeds to the right. H A + B
A − + H B+
stronger
acid
stronger
base
weaker
base
weaker
acid
A larger forward arrow means that products are favored.

29. 9. 3. Acid and Base Strength B
9.3 Acid and Base Strength B. Using Acid Strength to Predict the Direction of Equilibrium
If an acid–base reaction would form the stronger acid and base, equilibrium favors the reactants and little product forms. H A + B
A − + H B+
weaker
acid
weaker
base
stronger
base
stronger
acid
A larger reverse arrow means that reactants are favored.

30. HOW TO Predict the Direction of Equilibrium in an
9.3 Acid and Base Strength
HOW TO Predict the Direction of Equilibrium in an
Acid–Base Reaction
Are the reactants or products favored in
the following acid–base reaction?
Example
Identify the acid in the reactants and the
conjugate acid in the products.
Step [1]

31. HOW TO Predict the Direction of Equilibrium in an
9.3 Acid and Base Strength
HOW TO Predict the Direction of Equilibrium in an
Acid–Base Reaction
Determine the relative strength of the acid
and the conjugate acid.
Step [2]
Equilibrium favors the formation of the
weaker acid.
Step [3]

32. 9.3 Acid and Base Strength A. Relating Acid and Base Strength

33. Are the reactants or products favored at equilibrium in each reaction?
HF(g) OH- (aq) F-(aq) H2O(aq)
NH4+(aq) Cl NH3 (aq) HCl (aq)
HCO H3O H2CO3 (aq) H2O(aq)

34. 9.4 Equilibrium and Acid Dissociation Constants
For the reaction where an acid (HA) dissolves in
water,
HA(g) H2O(l)
H3O+(aq) (aq)
A −
the following equilibrium constant can be written:

35. 9.4 Equilibrium and Acid Dissociation Constants
Multiplying both sides by [H2O] forms a new constant, called the acid dissociation constant, Ka.

36. 9.4 Equilibrium and Acid Dissociation Constants

37. 9.5 Dissociation of Water
Water can behave as both a Brønsted–Lowry acid
and a Brønsted–Lowry base. Thus, two water molecules can react together in an acid–base reaction:

38. 9.5 Dissociation of Water
From the reaction of two water molecules, the following equilibrium constant expression can be written:
Multiplying both sides by [H2O]2 yields Kw, the ion-product constant for water.

39. 9.5 Dissociation of Water
Experimentally it can be shown that:

40. To calculate [OH−] when To calculate [H3O+] when
9.5 Dissociation of Water
To calculate [OH−] when
[H3O+] is known:
To calculate [H3O+] when
[OH−] is known:

41. 9.5 Dissociation of Water
If the [H3O+] in a cup of coffee is 1.0 x 10−5 M, then
the [OH−] can be calculated as follows:

42. If [H3O+] in blood is 4. 0 x 10-8 M, what is the value of [OH-]
If [H3O+] in blood is 4.0 x 10-8 M, what is the value of [OH-]? Is blood acidic, basic or neutral

43. Calculate the value of [H3O+] from the given [OH-] in each solution and label the solution as acidic or basic
[OH-] = 1x10-6 M
[OH-] = 5.2x10-11 M
[OH-] = 7.3x10-4 M

45. 9.6 The pH Scale A. Calculating pH
The lower the pH, the higher the concentration of
H3O+:

46. 9.6 The pH Scale A. Calculating pH

47. 9.6 The pH Scale B. Calculating pH Using a Calculator
Example:
If [H3O+] = 1.2 x 10–5 M for a solution, what is its pH?

48. 9.6 The pH Scale B. Calculating pH Using a Calculator
If the pH of a solution is 8.50, what is the [H3O+]?

49. 9.6 The pH Scale

50. What [H3O+] corresponds to each pH value?

51. 9. 7. Common Acid–Base Reactions A
9.7 Common Acid–Base Reactions A. Reaction of Acids with Hydroxide Bases
Neutralization reaction: An acid-base reaction that
produces a salt and water as products.

52. 9.7 Common Acid–Base Reactions
HOW TO Draw a Balanced Equation for a Neutralization
Reaction Between HA and MOH
Write a balanced equation for the reaction
of Mg(OH)2 with HCl.
Example
Identify the acid and base in the reactants
and draw H2O as one product.
Step [1]

53. 9.7 Common Acid–Base Reactions
HOW TO Draw a Balanced Equation for a Neutralization
Reaction between HA and MOH
Step [2]
Determine the structure of the salt.
The salt is formed from the parts of the acid and
base that are not used to form H2O.

54. 9.7 Common Acid–Base Reactions
HOW TO Draw a Balanced Equation for a Neutralization
Reaction between HA and MOH
Step [3]
Balance the equation.

55. Write a balanced equation for each acid-base reaction
HNO3(aq) + NaOH (aq) 
H2SO4(aq) + KOH(aq) 

56. 9. 7. Common Acid–Base Reactions A
9.7 Common Acid–Base Reactions A. Reaction of Acids with Hydroxide Bases
A net ionic equation contains only the species involved in a reaction.
HCl(aq) + NaOH(aq)
H—OH(l) + NaCl(aq)
Written as individual ions:
Omit the spectator ions, Na+ and Cl–.
What remains is the net ionic equation:

57. 9. 7. Common Acid–Base Reactions B
9.7 Common Acid–Base Reactions B. Reaction of Acids with Bicarbonate Bases
A bicarbonate base, HCO3−, reacts with one H+ to form carbonic acid, H2CO3.
H+(aq) + HCO3−(aq)
H2CO3(aq)
H2O(l) + CO2(g)
Carbonic acid then decomposes into H2O and CO2.
For example:
HCl(aq) + NaHCO3(aq)
NaCl(aq) + H2CO3(aq)
H2O(l) + CO2(g)

58. 9. 7. Common Acid–Base Reactions B
9.7 Common Acid–Base Reactions B. Reaction of Acids with Bicarbonate Bases
A carbonate base, CO32–, reacts with two H+ to form carbonic acid, H2CO3.
2 H+(aq) + CO32–(aq)
H2CO3(aq)
H2O(l) + CO2(g)
For example:
2 HCl(aq) + Na2CO3(aq)
2 NaCl(aq) + H2CO3(aq)
H2O(l) + CO2(g)

59. HBr + KOH 
HNO3 + Ca(OH)2
HCl + NaHCO3
H2SO4 + Mg(OH)2

60. HNO3 + LiOH 
H2SO4 + NaOH 
K2CO3 + HCl 
HI + NaHCO3

61. 9.8 The Acidity and Basicity of Salt Solutions
A salt can form an acidic, basic, or neutral
solution depending on whether its cation and anion
are derived from a strong or weak acid and base.
For the salt M+A−:
The cation M+ comes from the base.
The anion A− comes from the acid.

62. 9.8 The Acidity and Basicity of Salt Solutions
NaCl
Na+
from NaOH
strong base
Cl−
from HCl
strong acid
A salt derived from a strong acid and strong base
forms a neutral solution (pH = 7).

63. 9.8 The Acidity and Basicity of Salt Solutions
NaHCO3
Na+
from NaOH
strong base
HCO3−
from H2CO3
weak acid
A salt derived from a strong base and a weak acid
forms a basic solution (pH > 7).

64. 9.8 The Acidity and Basicity of Salt Solutions
NH4Cl
NH4+
from NH3
weak base
Cl−
from HCl
strong acid
A salt derived from a weak base and a strong acid
forms an acidic solution (pH < 7).

65. 9.8 The Acidity and Basicity of Salt Solutions
Thus, the ion derived from the stronger acid or
base determines whether the solution is acidic
or basic.

66. Determine whether each salt forms an acidic, basic, or neutral solution.KI
K2CO3
Ca(NO3)2

67. 9.9 Titration
To determine the concentration of an acid or base in a solution, we carry out a titration.
If we want to know the concentration of an acid solution, a base of known concentration is added slowly until the acid is neutralized.
When the acid is neutralized:
# of moles of acid = # of moles of base
This is called the end point of the titration.

68. 9.9 Titration

69. M (mol/L) conversion factor
9.9 Titration
Determining an unknown molarity from titration data
requires three operations:
mole–mole
conversion
factor
[2]
Moles of
base
Moles of
acid
M (mol/L) conversion factor
M (mol/L)
conversion
factor
[1]
[3]
Volume of base
Volume of
acid

70. HOW TO Determine the Molarity of an Acid Solution
9.9 Titration
HOW TO Determine the Molarity of an Acid Solution
from Titration
What is the molarity of an HCl solution if
22.5 mL of a M NaOH solution are
needed to titrate a 25.0 mL sample of the
acid?
Example

71. HOW TO Determine the Molarity of an Acid Solution
9.9 Titration
HOW TO Determine the Molarity of an Acid Solution
from Titration
Step [1]
Determine the number of moles of base
used to neutralize the acid.

72. HOW TO Determine the Molarity of an Acid Solution
9.9 Titration
HOW TO Determine the Molarity of an Acid Solution
from Titration
Determine the number of moles of acid
that react from the balanced chemical
equation.
Step [2]

73. HOW TO Determine the Molarity of an Acid Solution
9.9 Titration
HOW TO Determine the Molarity of an Acid Solution
from Titration
Step [3]
Determine the molarity of the acid from
the number of moles and the known
volume.

74. What is the molarity of an HCl solution if 25. 5 mL of a 0
What is the molarity of an HCl solution if 25.5 mL of a 0.24 M NaOH solution are needed to neutralize 15.0 mL of the sample?

75. What is the molarity of an HCl solution if 17. 2 mL of 0
What is the molarity of an HCl solution if 17.2 mL of 0.15 M NaOH are needed to neutralize 5.00 mL of the sample?