Mec chapter 8
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Transcript of Mec chapter 8
Chapter 8
Acids and Bases and Oxidation-Reduction
Denniston Topping Caret
7th Edition
Copyright The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
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 bases1. Arrhenius Theory
2. Brønsted-Lowry Theory
8.1
Aci
ds a
nd B
ases
• 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)
Arrhenius Theory of Acids and Bases
8.1
Aci
ds a
nd B
ases
Arrhenius Theory of 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)
8.1
Aci
ds a
nd B
ases
Arrhenius Theory of 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
8.1
Aci
ds a
nd B
ases
Brønsted-Lowry Theory of 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
HCl(aq) + H2O(l) Cl-(aq) + H3O+(aq)
What donated the proton? HClIs it an acid or base? Acid
What accepted the proton? H2OIs it an acid or base? Base
Brønsted-Lowry Theory of Acids and Bases
base
8.1
Aci
ds a
nd B
ases
acid
.
base acidNH3(aq) + H2O(l) NH4
+(aq) + OH-(aq)
8.1
Aci
ds a
nd B
ases
Brønsted-Lowry Theory of Acids and Bases
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
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
Aci
ds a
nd B
ases
8.1
Aci
ds a
nd B
ases
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
Aci
ds a
nd B
ases
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
CH3COOH(aq) + H2O(l) CH3COO-(aq) + H3O+(aq)
H2CO3(aq) + H2O(l) HCO3-(aq) + H3O+(aq)
8.1
Aci
ds a
nd B
ases
Weak Acids
• Weak acids and bases – only a small percent dissociates (Weak electrolytes)
• Weak acid examples:– Acetic acid:
– Carbonic Acid:
• Weak base examples:– Ammonia:
– Pyridine:
– Aniline:C6H5NH2(aq) + H2O(l) C6H5NH3
+(aq) + OH-(aq)
C5H5NH2(aq) + H2O(l) C5H5NH3+(aq) + OH-(aq)
NH3(aq) + H2O(l) NH4+(aq) + OH-(aq)
8.1
Aci
ds a
nd B
ases
Weak Bases
• The acid base reaction can be written in the general form:
• Notice the reversible arrows
• The products are also an acid and base called the conjugate acid and base
acid baseHA + B A– + HB+
8.1
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ds a
nd B
ases
Conjugate Acids and Bases
acid base
• 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
base acid
HA + B A- + HB+
8.1
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ds a
nd B
ases
HA + B A– + HB+
8.1
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ds a
nd B
ases
Acid-Base Dissociation
• 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
Aci
ds a
nd B
ases
Conjugate Acid-Base Pairs
• Which acid is stronger:
HF or HCN? HF
• Which base is stronger:
CN- or H2O? CN -
8.1
Aci
ds a
nd B
ases
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
• Pure water is virtually 100% molecular
• Very small number of molecules dissociate– Dissociation of acids and bases is often called
ionization
• Called autoionization
• Very weak electrolyte
H2O(l) + H2O(l) H3O+(aq) + OH-(aq)
8.1
Aci
ds a
nd B
ases
The Dissociation of Water
• 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:
Remember, liquids are not included in equilibrium expressions
]OH][O[HK -3eq
H2O(l) + H2O(l) H3O+(aq) + OH-(aq)
8.1
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ds a
nd B
ases
Hydronium Ion
• This constant is called the ion product for water and has the symbol Kw
• Since [H3O+] = [OH-] = 1.0 x 10-7 M, what is the value for Kw?
– 1.0 x 10-14
– It is unitless
]OH][O[HK -3w
8.1
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ds a
nd B
ases
Ion Product of Water
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+]
• 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-]8.2
pH: A
Mea
sure
men
t S
cale
for
Aci
ds a
nd B
ases A Definition of 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 pH8.
2 pH
: A M
easu
rem
ent
Sca
le f
or A
cids
and
Bas
es Measuring 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:
8.2
pH: A
Mea
sure
men
t S
cale
for
Aci
ds a
nd B
ases Calculating pH
pH = -log[H3O+]
1 x 10-14 = [H3O+][OH-]
8.2
pH: A
Mea
sure
men
t S
cale
for
Aci
ds a
nd B
ases
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
= -log [H3O+]
= -log [1.0 x 10-4]
= -[-4.00] = 4.00
pH = -log[H3O+]
8.2
pH: A
Mea
sure
men
t S
cale
for
Aci
ds a
nd B
ases Calculating [H3O+] From pH
What is the [H3O+] of a solution with pH = 6.00?
• 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+]
pH = -log[H3O+]
8.2
pH: A
Mea
sure
men
t S
cale
for
Aci
ds a
nd B
ases 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
• 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]
= 11.00
1 x 10-14 = [H3O+][OH-]
pH = -log[H3O+]
8.2
pH: A
Mea
sure
men
t S
cale
for
Aci
ds a
nd B
ases
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 [2.5 x 10-4]
• = 3.60
pH = -log[H3O+]
8.2
pH: A
Mea
sure
men
t S
cale
for
Aci
ds a
nd B
ases Calculating [OH-] From pH
What is the [OH-] of a solution with pH = 4.95?
• First find [H3O+] • 4.95 = -log [H3O+]
• [H3O+] = 10-4.95 • [H3O+] = 1.12 x 10-5
• Now solve for [OH-]• [OH-] = 1 x 10-14 / 1.12 x 10-5
= 1.0 x 10-9
pH = -log[H3O+]
1 x 10-14 = [H3O+][OH-]
The pH Scale8.
2 pH
: A M
easu
rem
ent
Sca
le f
or A
cids
and
Bas
es
1.0 x 100 0.001.0 x 10-1 1.001.0 x 10-2 2.001.0 x 10-3 3.001.0 x 10-4 4.001.0 x 10-5 5.001.0 x 10-6 6.001.0 x 10-7 7.00
For a strong acidHCl molarity pH
Mor
e A
cidi
c
1.0 x 100 14.001.0 x 10-1 13.001.0 x 10-2 12.001.0 x 10-3 11.001.0 x 10-4 10.001.0 x 10-5 9.001.0 x 10-6 8.001.0 x 10-7 7.00
For a strong baseNaOH molarity pH
Mor
e ba
sic
Each 10 fold change in concentration changes the pH by one unit
8.2
pH: A
Mea
sure
men
t S
cale
for
Aci
ds a
nd B
ases
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
8.3
Rea
ctio
ns B
etw
een
Aci
ds a
nd B
ases
• The net ionic neutralization reaction is more 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
Net Ionic Neutralization Reaction
Buret – long glass tube calibrated in mL which contains the standard solution
Flask contains a solution of unknown concentration plus indicator
Indicator – a substance which changes color as pH changes
Standard solution is slowly added until the color changes
The equivalence point is when the moles of H3O+ and OH- are equal
8.3
Rea
ctio
ns B
etw
een
Aci
ds a
nd B
ases
Acid – Base Titration
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
Aci
d-B
ase
Buf
fers
• 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
CH3COOH(aq) + H2O(l) CH3COO-(aq) + H3O+(aq)
The Buffer Process
CH3COOH(aq) + H2O(l) CH3COO-(aq) + H3O+(aq)8.4
Aci
d-B
ase
Buf
fers
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+
CH3COOH(aq) + H2O(l) CH3COO-(aq) + H3O+(aq)8.4
Aci
d-B
ase
Buf
fers
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
Aci
d-B
ase
Buf
fers
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