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Chapter 8
Acids and Bases and Oxidation-Reduction
Denniston Topping Caret
6th 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
• Acid - a substance, when dissolved in water, dissociates to produce hydrogen ionsHydrogen ion: H+ also called “protons”
HCl is an acid:
HCl(aq) H+(aq) + Cl-(aq)
Arrhenius Theory of Acids and Bases
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)
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
Brønsted-Lowry Theory of Acids and Bases
• Acid - proton donor
• Base - proton acceptoroNotice that acid and base are not defined
using wateroWhen 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
baseacid
.
base acidNH3(aq) + H2O(l) NH4
+(aq) + OH-(aq)
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 propertiesoAmphiprotic – a substance possessing both acid
and base propertiesoWater is the most commonly used solvent for
both acids and baseso Solute-solvent interactions between water and
both acids and bases promote solubility and dissociation
Acid and Base Strength
• Acid and base strength – degree of dissociationoNot a measure of concentrationo Strong acids and bases – reaction with water is
virtually 100% (Strong electrolytes)
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)
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)
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+
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+
HA + B A- + HB+
Acid-Base Dissociation
• The reversible arrow isn’t always writteno Some acids or bases essentially dissociate 100%
oOne way arrow is used
• HCl + H2O Cl- + H3O+ oAll of the HCl is converted to Cl-
oHCl is called a strong acid – an acid that dissociates 100%
• Weak acid - one which does not dissociate 100%
Conjugate Acid-Base Pairs
• Which acid is stronger:
HF or HCN? HF
• Which base is stronger:
CN- or H2O? CN -
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 dissociateoDissociation of acids and bases is often called
ionization
• Called autoionization
• Very weak electrolyte
H2O(l) + H2O(l) H3O+(aq) + OH-(aq)
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)
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?
o 1.0 x 10-14
o It is unitless
]OH][O[HK -3w
+=
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 solutiono 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 pHo if know one concentration, can calculate the
other
o if add an acid, [H3O+] and [OH-]
o if add a base, [OH-] and [H3O+]
o [H3O+] = [OH-] when equal amounts of acid and base are present
• In each of these cases 1 x 10-14 = [H3O+][OH-]
A Definition of pH
• pH of a solution can be:o Calculated if the concentration of either is
known• [H3O+] • [OH-]
oApproximated using indicator / pH paper that develops a color related to the solution pH
oMeasured using a pH meter whose sensor measures an electrical property of the solution that is proportional to pH
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:
Calculating pH
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?
oHCl is a strong acid and dissociates in watero If 1 mol HCl is placed in 1 L of aqueous
solution it produces 1 mol [H3O+]o 1.0 x 10-4 M HCl solution has [H3O+]=1.0x10-4M
= -log [H3O+]
= -log [1.0x10-4]
= -[-4.00] = 4.00
pH = -log[H3O+]
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+]
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+]
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]
• = 3.6
pH = -log[H3O+]
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+] = 1.122 x 10-5
• Now solve for [OH-]• [OH-] = 1 x 10-14 / 1.122 x 10-5
= 8.91 x 10-10
pH = -log[H3O+]
1 x 10-14 = [H3O+][OH-]
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
o Show only the changed componentso Omit any ions appearing the same on both sides of
equation = Spectator Ions
H+ + OH- H2O
• 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
Acid – Base Titration
Determine the Concentration of a Solution of Hydrochloric Acid• Place a known volume of acid whose concentration
is not known into a flask• Add an indicator, experience guides selection, here
phenol red is good• Known concentration of NaOH is placed in a buret• Drip NaOH into the flask until the indicator
changes color
Determine the Concentration of a Solution of Hydrochloric Acid
• Indicator changes color o equivalence point is reached
o mole OH- = mole H3O+ present in the unknown acid
• Volume dispensed from buret is determined• Calculate acid concentration:
o Volume of Hydrochloric Acid: 25.00 mLo Volume of NaOH added: 35.00 mLo Concentration of NaOH: 0.1000 Mo Balanced reaction shows that HCl and NaOH react 1:1
Determine the Concentration of a Solution of Hydrochloric Acid
• 35.00 mL NaOH x 1L NaOH x 0.1000 mol NaOH 103 mL NaOH 1L NaOH
= 3.500 x 10-3 mol NaOH• 3.500 x 10-3 mol NaOH x 1 mol HCl 1 mol NaOH
= 3.500 x 10-3 mol HCl– this amount of HCl is contained in 25.00 mL
• 3.500 x 10-3 mol HCl x 103 mL HCl 25.00 mL HCl 1 L HCl• = 1.400 x 10-1 mol HCl / L HCl = 0.1400 M HCl
• The previous examples have the acid and base at a 1:1 combining ratioo Not all acid-base pairs do this
• Polyprotic substance – donates or accepts more than one proton per formula unito Hydrochloric acid is monoprotic, producing one H+ ion for each unit of HCl
o Sulfuric acid is diprotic, each unit of H2SO4 produces 2 H+ ions
H2SO4(aq) + 2NaOH(aq) Na2SO4(aq) + 2 H2O(l)
Polyprotic Substances
Step 1.
H2SO4(aq) + H2O(l) HSO4-(aq) + H3O+(aq)
Step 2.
HSO4-(aq) + H2O(l) SO4
2-(aq) + H3O+(aq)
• In Step 1 H2SO4 behaves as a strong acid – dissociating completely
• In Step 2 HSO4-( behaves as a weak acid – reversibly dissociating,
note the double arrow
Dissociation of Polyprotic Substances
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
• Buffers act to establish an equilibrium between a conjugate acid – base pair
• Buffers consist of eithero a weak acid and its salt (conjugate base)o a weak base and its salt (conjugate acid)
o Acetic acid (CH3COOH) with sodium acetate (CH3COONa)
• An equilibrium is established in solution between the acid and the salt anion• A buffer is LeChatelier’s Principle in action
CH3COOH(aq) + H2O(l) CH3COO-(aq) + H3O+(aq)
The Buffer Process
CH3COOH(aq) + H2O(l) CH3COO-(aq) + H3O+(aq)
Addition of Base (OH-) to a Buffer Solution
• Adding a basic substance to a buffer causes changeso The OH- will react with the H3O+ producing watero Acid in the buffer system dissociates to replace
the H3O+ consumed by the added baseo 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)
Addition of Acid (H3O+) to a Buffer Solution
• Adding an acidic substance to a buffer causes changeso The H3O+ from the acid will increase the overall
H3O+ o Conjugate base in the buffer system reacts with the
H3O+ to form more acido 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
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
• Buffering process is an equilibrium reaction described by an equilibrium-constant expression
In acids, this constant is Ka
• If you want to know the pH of the buffer, solve for [H3O+], then calculate pH
COOH]CH[
]COOCH][OH[K
3
-33
a
+
=
Preparation of a Buffer Solution
CH3COOH(aq) + H2O(l) CH3COO-(aq) + H3O+(aq)
Henderson-Hasselbach Equation
• Solution of equilibrium-constant expression and pH can be combined into one operation
• Henderson-Hasselbach Equation is this combined expression
• Using these two equations:
o pKa = -log Ka just as pH = -log[H3O+]
o pKa = pH – log ( [CH3COO-] / [CH3COOH] )
o Henderson-Hasselbach –
o pH = pKa + log( [CH3COO-] / [CH3COOH] )
CH3COOH(aq) + H2O(l) CH3COO-(aq) + H3O+(aq)
COOH]CH[
]COOCH][OH[K
3
-33
a
+
=
Henderson-Hasselbach Equation
• pH = pKa + log( [CH3COO-] / [CH3COOH] ) can be rewritten
pH = pKa + log ( [conjugate base] / [weak acid])
8.5 Oxidation-Reduction Processes
• Oxidation-reduction processes are responsible for many types of chemical change
• Oxidation - defined by one of the following o loss of electrons o loss of hydrogen atomso gain of oxygen atoms
• Example: NaNa+ + e-
o Oxidation half reaction
• Reduction - defined by one of the following:o gain of electronso gain of hydrogeno loss of oxygen
• Example: Cl + e- Cl-
o Reduction half reaction
• Cannot have oxidation without reduction.
Oxidation-Reduction Processes
Na + Cl Na+ + Cl-
Oxidizing Agent• Is reduced• Gains electrons• Causes oxidation
Reducing Agent• Is oxidized• Loses electrons• Causes reduction
Oxidation and Reduction as Complementary Processes
Na Na+ + e-
Cl + e- Cl-
Applications of Oxidation and Reduction
• Corrosion - the deterioration of metals caused by an oxidation-reduction process
Example: rust (oxidation of iron)
4Fe(s) + 3O2(g) 2Fe2O3(s)
• Combustion of Fossil Fuels
Example: natural gas furnaces
CH4(g) + 2O2(g) CO2(g) + 2H2O(g)
• Bleaching
• Most bleaching agents are oxidizing agents
• The oxidation of the stains produces compounds that do not have color
Example: Chlorine bleach - sodium hypochlorite (NaOCl)
Applications of Oxidation and Reduction
Biological Processes
• Respirationo Electron-transport chain of aerobic
respiration uses reversible oxidation and reduction of iron atoms in cytochrome c
• Metabolismo Break down of molecules into smaller pieces
by enyzmes
• Is Zn oxidized or reduced?• Oxidized
• Copper is reduced
Voltaic Cells
• Voltaic cell – electrochemical cell that converts stored chemical energy into electrical energy
• Let’s consider the following reaction:
Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s)
Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s)
Voltaic Cells
• If the two reactants are placed in the same flask they cannot produce electrical current
• A voltaic cell separates the two half reactions
• This makes the electrons flow through a wire to allow the oxidation and reduction to occur
Zn Zn2+ + 2e-
Oxidationanode – electrode
where oxidation occurs
Cu2+ + 2e- CuReduction
cathode – electrode where reduction occurs
Voltaic Cell Generating Electrical Current
Electrolysis
• Electrolysis reactions – uses electrical energy to cause nonspontaneous oxidation-reduction reactions to occur
• These reactions are the reverse of a voltaic cell
Rechargeable battery• When powering a device behaves as voltaic cell• With time the chemical reaction nears
completion• Battery appears to “run down”• Cell reaction is reversible when battery attached
to charger