Mec chapter 8

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Chapter 8
Acids and Bases and
Oxidation-Reduction

Denniston
Topping
Caret
7th
Edition

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

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)

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)

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

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

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

.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

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

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)

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

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)

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)

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

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

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%

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 -

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

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

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

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

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+]

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-]

Measuring pH
• 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

Calculating pH

• 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-]

Calculating pH from
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

Calculating [H3O+] From pH
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+]

Calculating the pH of a Base
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

Calculating pH from Acid
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

Calculating [OH-] From pH

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

The pH Scale

For a strong acid For a strong base

HCl molarity pH
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

The Importance of pH and
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

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

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

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

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

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)
– 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

Addition of Base (OH-) to a
Buffer Solution

• 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+

Addition of Acid (H3O+) to a
Buffer Solution

• 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

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

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