1. Acid-Base Equilibria: The Nature of Acids and Bases
What makes an Acid an Acid?
An acid possess a sour taste
An acid dissolves active metals like magnesium
An acid causes certain vegetable dyes to turn
characteristic colors
What makes a Base a Base?
A bases possess a bitter taste
A base feels slippery to the touch
A base causes certain vegetable dues to turn a
characteristic color

2. Acid-Base Equilibria: The Nature of Acids and Bases
The Arrhenius Definition of an Acid and a Base:
An acid is a substance that produces H+ ions in water
solutions
A base is a substance that produces OH- ions in a water
solution
If a solution contains more OH- ions than H+ ions we say that the solution is basic
If a solution contains more H+ ions than OH- ions we say that the solution is acidic

3. Acid-Base Equilibria: The Proton in Water
When HCl dissolves in water we write:
H2O
HCl(g)  H+(aq) + Cl-(aq) H : H+  O : H
hydronium ion (H3O+)

4. Acid-Base Equilibria: The Proton in Water
When HCl dissolves in water we write:
H2O
HCl(g)  H+(aq) + Cl-(aq)

5. Acid-Base Equilibria: The Proton in Water
When HCl dissolves in water we write:
H2O
HCl(g)  H+(aq) + Cl-(aq)
Acidic solutions are formed by a chemical reaction in which
and acid transfers a proton (H+) to water.
HCl(aq) + H2O(aq)  H3O+(aq) + Cl-(aq)

6. Acid-Base Equilibria: The Proton in Water
The Bronsted-Lowry Concept of Acids and Bases
An acid may be defined as a substance that is capable of
donating protons and a base may be defined as substance
than accepts protons. + H H : _ : : : Cl H + : : N H  Cl + H N H : : H H
Bronsted
acid
Bronsted
base

7. Acid-Base Equilibria: The Proton in Water
The Bronsted-Lowry Concept of Acids and Bases
Describing the relationship between the Arrhenius and Bronsted
definitions of acids and bases,
NH3(aq) + H2O(aq)  NH4+(aq) + OH-(aq)
The Bronsted definition says that H2O
is a an acid because it donated a proton
and NH3 is a base because it accepted
the proton
The Arrhenius definition says that this
solution is basic
Water is an acid

8. Acid-Base Equilibria: The Proton in Water
The Bronsted-Lowry Concept of Acids and Bases
HC2H3O2(aq) + H2O(aq)  H3O+(aq) + C2H3O2(aq)
The Bronsted definition says that acetic acid is a an acid because it donates a proton to water and and water is a base because it accepts the proton from the acetic acid
Water is also a
base

9. Acid-Base Equilibria: The Proton in Water
The Bronsted-Lowry Concept of Acids and Bases
HC2H3O2(aq) + H2O(aq)  H3O+(aq) + C2H3O2(aq)
Bronsted
acid
Bronsted
base
Conjugate
acid
Conjugate
base
NH3(aq) + H2O(aq)  NH4+(aq) + OH-(aq)
Bronsted
base
Bronsted
acid
Conjugate
acid
Conjugate
base
The stronger the acid, the weaker its conjugate base; the weaker
the acid, the stronger its conjugate base.

10. Acid-Base Equilibria: The Proton in Water
The Bronsted-Lowry Concept of Acids and Bases
Of course all this Bronsted-Lowry ‘stuff’ raises a number of questions.
If water can be an acid and a base, can it act as a proton donor and acceptor
to itself?
What makes one acid or base strong and another acid or base weak?

11. Acid-Base Equilibria: The Proton in Water
The Relative Strengths of Acids and Bases
What makes one acid or base strong and another acid or base weak?

12. Acid-Base Equilibria: The Proton in Water
The Dissociation of Water and the pH Scale
If water can be an acid and a base, can it act as a proton donor and acceptor
to itself?
Water is capable of auto-ionizing.
H2O (l)  H+ (aq) + OH- (aq) H O : +  ..
The reaction occurs to a very small extent; about 1 in 108 molecules is ionized
at any given moment
Protons transfer from one molecule to another at a rate of about 1000 times
per second

13. Acid-Base Equilibria: The Proton in Water
The Dissociation of Water and the pH Scale
If this is true: H2O (l)  H+ (aq) + OH- (aq)
[H+] [OH-]
Than this is true:
K =
[H2O]
And since water is a liquid and its concentration is therefore constant, this
expression may be written as:
Kw =
[H+] [OH-] where Kw is the ion product constant
and is equal to 1.0 x 10-14
Note that [H+] = [OH-] = 1.0 x 10-7 M, water is therefore said to be neutral
However, in most solutions these concentrations vary. If
If [H+] > [OH-] , solution is acidic
If [H+] > [OH-] , solution is basic

14. Acid-Base Equilibria: The Proton in Water
The Bronsted-Lowry Concept of Acids and Bases
Sample exercise: Indicate whether each of the following solutions is neutral, acidic,
or basic: (a) [H+] = 2 x 10-5 M, (b) [OH-] = M, ( c) [OH-] = 1.0 x 10-7 M
Sample exercise: Calculate the concentration of H+ (aq) in (a) a solution in which the
[OH-] is 0.020M, (b) a solution in which the [OH-] = 2.5 x 10-6 M. Indicate whether
the solution is acidic or basic

15. Acid-Base Equilibria: The Proton in Water
The Dissociation of Water and the pH Scale

16. Acid-Base Equilibria: The Proton in Water
The Dissociation of Water and the pH Scale
Because the concentration of H+ ions is often quite small, it can be conveniently expressed in terms of pH = -log [H+]
For example, solution with a [H+] = 2. 5 x 105 has a pH of:
pH = -log [2. 5 x 10-5 ] = 4.6
Likewise a solution with a pH of 3.8 has a H+ concentration of:
Antilog -3.8= 1.58 x 104 M

17. Acid-Base Equilibria: The Proton in Water
The Dissociation of Water and the pH Scale
In a sample of lemon juice, [H+] = 3.8 x 10-4 M. What is the pH. A commonly
available window cleaner has a [H+] = 5.3 x 10-9 M
In a sample of freshly pressed apple juice has a pH of Calculate the [H]
Now, you
try it!

18. Sooooo… if Kw = [H+] [OH-] and Kw = 1.0 x 10-14, then:
Acid-Base Equilibria: The Proton in Water
The Dissociation of Water and the pH Scale
Because the concentration of OH- ions is often quite small, it can be conveniently expressed in terms of p)H = -log [OH-]
Now lets think about this, if the [H] = [OH-], then the pH = pOH = 7
Sooooo… if Kw = [H+] [OH-] and Kw = 1.0 x 10-14, then:
-log Kw = 14 = pH + pOH
Sample exercise: What is the pH of a solution with a pOH of 2.5? Is the solution
acidic or basic?

19. Acid-Base Equilibria: The Proton in Water
Measuring the pH Using Indicators

20. Acid-Base Equilibria: The Differences Between Strong and Weak Acids
HX H X-
HX HX H X-
Initial Equilibrium Initial Equilibrium

21. Acid-Base Equilibria: The Differences Between Strong and Weak Acids

22. Acid-Base Equilibria: The Differences Between Strong and Weak Acids
Dealing with a Strong Acid:
What is the pH of M solution of HCl?
Dealing with a weak acid that is only partially ionizable:
[H+][X-]
Since HX (aq)  H+(aq) + X-(aq), then Ka =
[HX]
The smaller the value of the acid dissociation constant Ka, the weaker the acid
What is the Ka of a 0.10 M solution of formic acid (HCHO2) which has a pH = 2.38?
HCHO2  H CHO2 I C E

23. Acid-Base Equilibria: The Differences Between Strong and Weak Acids
What is the concentration of H+ ions in a 0.10 M solution of HC2H3O2 (Ka = 1.8 x 10-5)
HC2H3O2  H+ + C2H3O2 I C E
What is the pH of the solution?
What is the percent ionization of this solution?

24. Acid-Base Equilibria: The Differences Between Strong and Weak Acids
What is the pH and percent ionization of a 0.20 M solution of HCN? Ka = 4.9 x I C E

25. Acid-Base Equilibria: Dealing with Polyprotic Acids
Substances that are capable of furnishing more than one proton to water are called
polyprotic acids.
H2SO3(aq)  H+(aq) + HSO3-(aq) K a1 = 1.7 x 10-2
HSO3-(aq)  H+(aq) + SO32-(aq) K a2 = 6.4 x 10-8
Because Ka1 is so much larger than subsequent dissociation constants for most
polyprotic acids, almost all the H+ (aq) in the solution come from the first
ionization reaction.

26. Acid-Base Equilibria: Dealing with Polyprotic Acids

27. Acid-Base Equilibria: Dealing with Polyprotic Acids
What is the pH of M solution of carbonic acid (H2CO3)
H2CO  H HCO3- I C E
HCO  H CO32- I C E

28. Acid-Base Equilibria: Strong Bases
The most common soluble strong Bases are the hydroxides of group IA and heavier
group 2A metals
What is the pH of a M solution of Ba(OH)2?

29. Acid-Base Equilibria: Dealing with Weak Bases
The base dissociation constant Kb refers to the equilibrium in which a base reacts
with H2O to from the conjugate acid and OH-
Weak base + H2O  conjugate acid + OH-
NH3 (aq) + H2O (l)  NH4(aq) + OH-(aq)
[NH4+] [OH-]
Kb =
[NH3]
Calculate the [OH-] in a 0.15 M solution of NH3.
NH H2O  NH OH- I C E

30. Acid-Base Equilibria: Classes of Weak Acids
Amines
Anions of Weak Acids

31. Acid-Base Equilibria: Anions of Weak Acids
HC2H3O2(aq) + H2O(aq)  H3O+(aq) + C2H3O2- (aq)
Bronsted
acid
Bronsted
base
Conjugate
acid
Conjugate
base
A second class of weak base is composed of the anions of weak acids Anions
of weak acids can be incorporated into salts
NaC2H3O2  Na+(aq) + C2H3O2- (aq)
C2H3O2- + H2O  HC2H3O2 + OH- Kb = 5.6 x 1010

32. Acid-Base Equilibria: Anions of Weak Acids
Calculate the pH of a 0.01 M solution of sodium hypochlorite (NaClO)
+ H2O  OH- I C E

33. Acid-Base Equilibria: Anions of Weak Acids
Now it’s you turn: the Kb for BrO- is 5.0 x Calculate the pH of a M
solution of NaBrO

34. Acid-Base Equilibria: Relationship Between Ka and Kb
NH4+(aq)  NH3(aq) + H+ (aq)
NH3(aq) + H2O NH4+(aq) + OH- (aq)
[H+][NH3]
[NH4][OH- ]
Ka =
Kb =
[NH4+]
[NH3]
NH4+(aq)  NH3(aq) + H+ (aq)
NH3(aq) + H2O(l) NH4+(aq) + OH- (aq)
H2O  H+(aq) + OH-(aq)
When two reactions are added to give a third reaction, the equilibrium constant
for the third reaction reaction is given by the product of the equilibrium constants
for the two added reactions
Ka x Kb = Kw
pKa + pKb = pKw

35. Acid-Base Equilibria: Relationship Between Ka and Kb
Calculate the (a) base-dissociation constant, Kb, for the fluoride ion, is the pKa of
HF = 3.17
pKa = -log Ka
3.17 = -log Ka
Antilog = 6.76 x 10-4
Since
Ka x Kb = Kw
(6.76 x 10-4)x Kb = 1.0 x 10-14
Kb = 1.0 x 10-14/ 6.76 x 10-4 = 1.5 x 10-11

36. Acid-Base Equilibria: Relationship Between Ka and Kb
Calculate the pKb for carbonic acid (Ka = 4.3 x 10-7)
Now it’s your turn

37. Acid-Base Equilibria: Acid-Base Properties of Salt Solutions
Anions of weak acids, HX, are basic and will react with H2O to produce OH-
X- (aq) + H2O (l)  HX(aq) + OH-(aq)
Anions of strong acids, such as NO3-, exhibit no basicitiy, these ions do not react
with water and consequently do not influence the pH
Anions of polyprotic acids, such as HCO3-, that still have ionizable protons
are capable of acting as either proton donors or acceptors depending upon the
magnitudes of the Ka or Kb
This last one requires
a bit of an explanation

38. Acid-Base Equilibria: Acid-Base Properties of Salt Solutions
Anions of polyprotic acids, such as HCO3-, that still have ionizable protons
are capable of acting as either proton donors or acceptors depending upon the
magnitudes of the Ka or Kb
Predict whether the salt Na2HPO4 will form an acidic or basic solution on dissolving
in water.
Na2HPO4  2Na+ (aq) + HPO4-
HPO4- acting like an acid
HPO4- (aq) + H2O  H3O+ + PO43-(aq)
K3 = 4.2 x 10-13
HPO4- acting like an base
HPO4- (aq) + H2O  H2PO42-(aq) + OH-(aq)
So HPO- is the conjugate base of H2PO4-. Since the K2 of H2PO4- = 6.2 x then: Kw
1.0 x 10-14
Kb = =
1.6 x 10-7 = Ka
6.2 x 10-8
Since Kb is larger than Ka, HPO4- will act like a base

39. Acid-Base Equilibria: Acid-Base Properties of Salt Solutions
Salt derived from a strong base and a strong acid will have a pH of 7
Salt derived from a strong base and a weak acid will have a pH above 7
Salt derived from a weak acid and a weak base depends upon whether
the dissolved ion acts as an acid or a base as determined by the size of
the Ka or Kb

40. Acid-Base Equilibria: Acid-Base Character and Chemical Structure
A substance HX will transfer a proton only if the H X bond, is already polarized
in the following way: H X
In ionic compounds such as NaH, the H atom possess a negative charge and behaves
like a proton acceptor.
Very strong bonds in an HX are less easily ionizable than weak bonds

41. Acid-Base Equilibria: Acid-Base Character and Chemical Structure