This document provides an overview of acids and bases including: 1) It describes the Arrhenius and Brønsted-Lowry theories of acids and bases, defining acids as proton donors and bases as proton acceptors. 2) It discusses strong vs. weak acids and bases based on their degree of dissociation in water, and introduces conjugate acid-base pairs. 3) It explains that water can act as both an acid and base, and discusses the autoionization of water and the definition of pH in terms of the hydronium ion concentration.

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Chapter 8
Acids and Bases and
Oxidation-Reduction
Denniston
Topping
Caret
7th
Edition

2. 8.1 Acids and Bases
• Acids: Taste sour, dissolve some metals,
cause plant dye to change color
• Bases: Taste bitter, are slippery, are
corrosive
• Two theories that help us to understand
the chemistry of acids and bases
1. Arrhenius Theory
2. Brønsted-Lowry Theory

3. Arrhenius Theory of Acids
and Bases
8.1 Acids and Bases
• Acid - a substance, when dissolved in
water, dissociates to produce hydrogen
ions
– Hydrogen ion: H+ also called “protons”
HCl is an acid:
HCl(aq) → H+(aq) + Cl-(aq)

4. Arrhenius Theory of Acids
and Bases
8.1 Acids and Bases
• Base - a substance, when dissolved in
water, dissociates to produce hydroxide
ions
NaOH is a base
NaOH(aq) → Na+(aq) + OH-(aq)

5. Arrhenius Theory of Acids
and Bases
8.1 Acids and Bases
• Where does NH3 fit?
• When it dissolves in water it has basic
properties but it does not have OH- ions in
it
• The next acid-base theory gives us a
broader view of acids and bases

6. Brønsted-Lowry Theory of
Acids and Bases
8.1 Acids and Bases
• Acid - proton donor
• Base - proton acceptor
– Notice that acid and base are not defined
using water
– When writing the reactions, both accepting
and donation are evident

7. Brønsted-Lowry Theory of
Acids and Bases
8.1 Acids and Bases
HCl(aq) + H2O(l) → Cl-(aq) + H3O+(aq)
acid base
What donated the proton? HCl
Is it an acid or base? Acid
What accepted the proton? H2O
Is it an acid or base? Base

8. .Brønsted-Lowry Theory of
Acids and Bases
8.1 Acids and Bases
NH3(aq) + H2O(l) NH4+(aq) + OH-(aq)
base acid
Now, let us look at NH3 and see why it is a
base.
Did NH3 donate or accept a proton? Accept
Is it an acid or base? Base
What is water in this reaction? Acid

9. 8.1 Acids and Bases Acid-Base Properties of Water
• Water possesses both acid and base
properties
– Amphiprotic - a substance possessing both acid
and base properties
– Water is the most commonly used solvent for
both acids and bases
– Solute-solvent interactions between water and
both acids and bases promote solubility and
dissociation

10. 8.1 Acids and Bases Acid and Base Strength
• Acid and base strength – degree of
dissociation
– Not a measure of concentration
– Strong acids and bases – reaction with water is
virtually 100% (Strong electrolytes)

11. 8.1 Acids and Bases Strong Acids and Bases
• Strong Acids:
– HCl, HBr, HI Hydrochloric Acid, etc.
– HNO3 Nitric Acid
– H2SO4 Sulfuric Acid
– HClO4 Perchloric Acid
• Strong Bases:
– NaOH, KOH, Ba(OH)2
– All metal hydroxides

12. 8.1 Acids and Bases Weak Acids
• Weak acids and bases – only a small
percent dissociates (Weak electrolytes)
• Weak acid examples:
– Acetic acid:
CH3COOH(aq) + H2O(l) CH3COO-(aq) + H3O+(aq)
– Carbonic Acid:
H2CO3(aq) + H2O(l) HCO3-(aq) + H3O+(aq)

13. 8.1 Acids and Bases Weak Bases
• Weak base examples:
– Ammonia:
NH3(aq) + H2O(l) NH4+(aq) + OH-(aq)
– Pyridine:
C5H5NH2(aq) + H2O(l) C5H5NH3+(aq) + OH-(aq)
– Aniline:
C6H5NH2(aq) + H2O(l) C6H5NH3+(aq) + OH-(aq)

14. 8.1 Acids and Bases Conjugate Acids and Bases
• The acid base reaction can be written in
the general form:
HA + B A– + HB+
acid base
• Notice the reversible arrows
• The products are also an acid and base
called the conjugate acid and base

15. 8.1 Acids and Bases
HA + B A- + HB+
acid base base acid
• Conjugate acid - what the base becomes
after it accepts a proton
• Conjugate base - what the acid becomes
after it donates its proton
• Conjugate acid-base pair - the acid and
base on the opposite sides of the
equation

16. 8.1 Acids and Bases Acid-Base Dissociation
HA + B A– + HB+
• The reversible arrow isn’t always written
– Some acids or bases essentially dissociate 100%
– One way arrow is used
• HCl + H2O → Cl- + H3O+
– All of the HCl is converted to Cl-
– HCl is called a strong acid – an acid that
dissociates 100%
• Weak acid - one which does not dissociate
100%

17. 8.1 Acids and Bases Conjugate Acid-Base Pairs
• Which acid is
stronger:
HF or HCN? HF
• Which base is
stronger:
CN- or H2O? CN -

18. 8.1 Acids and Bases Acid-Base Practice
• Write the chemical reaction for the following
acids or bases in water
• Identify the conjugate acid-base pairs
1. HF (a weak acid)
2. H2S (a weak acid)
3. HNO3 (a strong acid)
4. CH3NH2 (a weak base)
Note: The degree of dissociation also defines weak
and strong bases

19. 8.1 Acids and Bases The Dissociation of Water
• Pure water is virtually 100% molecular
• Very small number of molecules dissociate
– Dissociation of acids and bases is often called
ionization
H2O(l) + H2O(l) H3O+(aq) + OH-(aq)
• Called autoionization
• Very weak electrolyte

20. 8.1 Acids and Bases Hydronium Ion
• H3O+ is called the hydronium ion
• In pure water at room temperature:
– [H3O+] = 1 x 10-7 M
– [OH-] = 1 x 10-7 M
• What is the equilibrium expression for:
H2O(l) + H2O(l) H3O+(aq) + OH-(aq)
+
K eq = [H 3O ][OH ] -
Remember, liquids are not included in equilibrium
expressions

21. 8.1 Acids and Bases Ion Product of Water
• This constant is called the ion product for
water and has the symbol Kw
+
K w = [H 3O ][OH ] -
• Since [H3O+] = [OH-] = 1.0 x 10-7 M, what is
the value for Kw?
– 1.0 x 10-14
– It is unitless

22. 8.2 pH: A Measurement Scale
for Acids and Bases
• pH scale - a scale that indicates the acidity
or basicity of a solution
– Ranges from 0 (very acidic) to 14 (very basic)
• The pH scale is rather similar to the
temperature scale assigning relative values
of hot and cold
• The pH of a solution is defined as:
pH = -log[H3O+]

23. A Definition of pH
Scale for Acids and Bases
8.2 pH: A Measurement
• Use these observations to develop a concept
of pH
– if know one concentration, can calculate the
other
– if add an acid, [H3O+] ↑ and [OH-] ↓
– if add a base, [OH-] ↑ and [H3O+] ↓
– [H3O+] = [OH-] when equal amounts of acid
and base are present
• In each of these cases 1 x 10-14 = [H3O+][OH-]

24. Measuring pH
Scale for Acids and Bases
8.2 pH: A Measurement
• pH of a solution can be:
– Calculated if the concentration of either is
known
• [H3O+]
• [OH-]
– Approximated using indicator / pH paper
that develops a color related to the solution
pH
– Measured using a pH meter whose sensor
measures an electrical property of the
solution that is proportional to pH

25. Calculating pH
Scale for Acids and Bases
8.2 pH: A Measurement
• How do we calculate the pH of a solution
when either the hydronium or hydroxide
ion concentration is known?
• How do we calculate the hydronium or
hydroxide ion concentration when the pH
is known?
• Use two facts:
pH = -log[H3O+]
1 x 10-14 = [H3O+][OH-]

26. Calculating pH from
Scale for Acids and Bases
8.2 pH: A Measurement
Acid Molarity
What is the pH of a 1.0 x 10-4 M HCl solution?
– HCl is a strong acid and dissociates in water
– If 1 mol HCl is placed in 1 L of aqueous
solution it produces 1 mol [H3O+]
– 1.0 x 10-4 M HCl solution has [H3O+]=1.0x10-4M
pH = -log[H3O+]
= -log [H3O+]
= -log [1.0 x 10-4]
= -[-4.00] = 4.00

27. Calculating [H3O+] From pH
Scale for Acids and Bases
8.2 pH: A Measurement
What is the [H3O+] of a solution with pH = 6.00?
pH = -log[H3O+]
• 4.00 = -log [H3O+]
• Multiply both sides of equation by –1
• -4.00 = log [H3O+]
• Take the antilog of both sides
• Antilog -4.00 = [H3O+]
• Antilog is the exponent of 10
• 1.0 x 10-4 M = [H3O+]

28. Calculating the pH of a Base
Scale for Acids and Bases
8.2 pH: A Measurement
What is the pH of a 1.0 x 10-3 M KOH solution?
• KOH is a strong base (as are any metal hydroxides)
• 1 mol KOH dissolved and dissociated in aqueous
solution produces 1 mol OH-
• 1.0 x 10-3 M KOH solution has [OH-] = 1.0 x 10-3 M
1 x 10-14 = [H3O+][OH-]
• Solve equation for [H3O+] = 1 x 10-14 / [OH-]
• [H3O+] = 1 x 10-14 / 1.0 x 10-3 = 1 x 10-11
• pH = -log [1 x 10-11]
pH = -log[H3O+]
= 11.00

29. Calculating pH from Acid
Scale for Acids and Bases
8.2 pH: A Measurement
Molarity
What is the pH of a 2.5 x 10-4 M HNO3 solution?
• We know that as a strong acid HNO3 dissociates
to produce 2.5 x 10-4 M [H3O+]
pH = -log[H3O+]
• pH = -log [2.5 x 10-4]
• = 3.60

30. Calculating [OH-] From pH
Scale for Acids and Bases
8.2 pH: A Measurement
What is the [OH-] of a solution with pH = 4.95?
• First find [H3O+] pH = -log[H O+] 3
• 4.95 = -log [H3O+]
• [H3O+] = 10-4.95
• [H3O+] = 1.12 x 10-5 1 x 10-14 = [H3O+][OH-]
• Now solve for [OH-]
• [OH-] = 1 x 10-14 / 1.12 x 10-5
= 1.0 x 10-9

31. 8.2 pH: A Measurement
Scale for Acids and Bases
The pH Scale

32. For a strong acid For a strong base
Scale for Acids and Bases
HCl molarity pH
8.2 pH: A Measurement
NaOH molarity pH
1.0 x 100 0.00 1.0 x 100 14.00
1.0 x 10-1 1.00 1.0 x 10-1 13.00
More basic
1.0 x 10-2 2.00 1.0 x 10-2 12.00
1.0 x 10-3 3.00 1.0 x 10-3 11.00
1.0 x 10-4 4.00 1.0 x 10-4 10.00
1.0 x 10-5 5.00 1.0 x 10-5 9.00
1.0 x 10-6 6.00 1.0 x 10-6 8.00
ci d c Aer o M
1.0 x 10-7 7.00 1.0 x 10-7 7.00
Each 10 fold change in concentration
i
changes the pH by one unit

33. The Importance of pH and
Scale for Acids and Bases
8.2 pH: A Measurement
pH Control
Any change that takes place in aqueous solution
generally has at least some pH dependence
– Agriculture - crops grow best in soil with proper
pH
– Physiology - blood pH shift of 1 pH is fatal
– Acid Rain - lowers pH of water in aquatic
systems causing problems for native fishes
– Municipal services - sewage treatment and water
purification require optimal pH
– Industry - many processes require strict pH
control for cost-effective production

34. 8.3 Reactions Between Acids
and Bases
• Neutralization reaction - the reaction of an acid
with a base to produce a salt and water
HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)
Acid Base Salt Water
• Break apart into ions:
H+ + Cl- + Na+ + OH- →Na+ + Cl- + H2O
• Net ionic equation
– Show only the changed components
– Omit any ions appearing the same on both sides of
equation = Spectator ions
H+ + OH- → H2O

35. Net Ionic Neutralization Reaction
8.3 Reactions Between
• The net ionic neutralization reaction is more
Acids and Bases
accurately written:
H3O+(aq) + OH-(aq) → 2H2O(l)
• This equation applies to any strong acid / strong
base neutralization reaction
• An analytical technique to determine the
concentration of an acid or base is titration
• Titration involves the addition of measured
amount of a standard solution to neutralize the
second, unknown solution
• Standard solution - solution of known
concentration

36. 8.3 Reactions Between Acid – Base Titration
Buret – long glass
Standard solution
tube calibrated in mL
is slowly added
Acids and Bases
which contains the
until the color
standard solution
changes
The equivalence Indicator – a
point is when the substance which
moles of H3O+ changes color as
pH changes
and OH- are equal
Flask contains a
solution of unknown
concentration plus
indicator

37. 8.4 Acid-Base Buffers
• Buffer solution - solution which resists large
changes in pH when either acids or bases are
added
• These solutions are frequently prepared in
laboratories to maintain optimum conditions
for chemical reactions
• Buffers are also used routinely in
commercial products to maintain optimum
conditions for product behavior

38. 8.4 Acid-Base Buffers The Buffer Process
• Buffers act to establish an equilibrium between a
conjugate acid – base pair
• Buffers consist of either
– a weak acid and its salt (conjugate base)
– a weak base and its salt (conjugate acid)
CH3COOH(aq) + H2O(l) CH3COO-(aq) + H3O+(aq)
– Acetic acid (CH3COOH) with sodium acetate
(CH3COONa)
• An equilibrium is established in solution
between the acid and the salt anion
• A buffer is Le Chatelier’s principle in action

39. Addition of Base (OH-) to a
Buffer Solution
8.4 Acid-Base Buffers
• Adding a basic substance to a buffer causes
changes
– The OH- will react with the H3O+ producing water
– Acid in the buffer system dissociates to replace
the H3O+ consumed by the added base
– Net result is to maintain the pH close to the initial
level
• The loss of H3O+ (the stress) is compensated
by the dissociation of the acid to produce
more H3O+
CH3COOH(aq) + H2O(l) CH3COO-(aq) + H3O+(aq)

40. Addition of Acid (H3O+) to a
Buffer Solution
8.4 Acid-Base Buffers
• Adding an acidic substance to a buffer causes
changes
– The H3O+ from the acid will increase the overall
H3O+
– Conjugate base in the buffer system reacts with
the H3O+ to form more acid
– Net result is to maintain the H3O+ concentration
and the pH close to the initial level
• The gain of H3O+ (the stress) is compensated
by the reaction of the conjugate base to
produce more acid
CH3COOH(aq) + H2O(l) CH3COO-(aq) + H3O+(aq)

41. 8.4 Acid-Base Buffers Buffer Capacity
• Buffer capacity - a measure of the ability
of a solution to resist large changes in
pH when a strong acid or strong base is
added
• Also described as the amount of strong
acid or strong base that a buffer can
neutralize without significantly changing
pH