Acids & Bases
CHAPTER 16
(& part of CHAPTER 17)
Chemistry: The Molecular Nature of Matter, 6th edition
By Jesperson, Brady, & Hyslop
CHAPTER 16: Acids & Bases
Learning Objectives:
 Define Brønsted-Lowry Acid/Base
 Define Lewis Acid/Base
 Evaluate the strength of acids/bases
 Strong vs weak acids/bases
 Periodic trends
 Conjugate acids/bases
 Identify likely compounds that will form acids and
bases from the periodic table
 Acidic metal ions
 Acid/Base equilibrium:
 pH, pOH
 Ka, Kb, pKa, pKb
 Kw of water
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CHAPTER 16: Acids & Bases
Lecture Road Map:
① Brønsted-Lowry Acids/Bases
② Trends in acid strength
③ Lewis Acids & Bases
④ Acidity of hydrated metal ions
⑤ Acid/Base equilibrium
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CHAPTER 16 Acids & Bases
Brønsted-Lowry
Acid/Base
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Arrhenius
Acid/Base
Definition
Acid produces H3O+ in water
Base gives OH–
Acid-base neutralization
– Acid and base combine to produce water and a salt.
e.g. HCl(aq) + NaOH(aq)  H2O + NaCl(aq)
H3O+(aq) + Cl–(aq) + Na+(aq) + OH–(aq)
 2H2O + Cl–(aq) + Na+(aq)
• Many reactions resemble this without forming H3O+ or OH–
in solution
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Arrhenius
Acid/Base
Definition
Gas Phase Acid/Base chemistry not covered
by Arrhenius definition
e.g. NH3(g) + HCl(g)  NH4Cl(s)
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BrønstedLowry
Definition
• Acid = proton donor
• Base = proton acceptor
• Allows for gas phase acid-base reactions
e.g. HCl + H2O  H3O+ + Cl–
– HCl = acid
• Donates H+
– Water = base
• Accepts H+
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BrønstedLowry
Conjugate Acid-Base Pair
• Species that differ by H+
e.g. HCl + H2O  H3O+ + Cl–
• HCl = acid
• Water = base
• H3O+
– Conjugate acid of H2O
• Cl–
– Conjugate base of HCl
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BrønstedLowry
Example: Formic Acid
• Formic acid (HCHO2) is a weak acid
• Must consider equilibrium
– HCHO2(aq) + H2O
CHO2–(aq) + H3O+(aq)
• Focus on forward reaction
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BrønstedLowry
Example: Formic Acid
Now consider reverse reaction:
• Hydronium ion transfers H+ to CHO2–
• Formate Ion is the Brønsted Base
conjugate pair
HCHO2 + H2O
acid
base
H3O+ + CHO2
acid
base
conjugate pair
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Group
Problem
• Identify the conjugate partner for each
conjugate base
conjugate acid
HCl
NH3
HC2H3O2
CN–
HF
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Group
Problem
Write a reaction that shows that HCO3– is a
Brønsted acid when reacted with OH–
Write a reaction that shows that HCO3– is a
Brønsted base when reacted with H3O+(aq)
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Group
Problem
In the following reaction, identify the acid/base
conjugate pairs.
(CH3)2NH + H2SO4 → (CH3)2NH+ + HSO4–
A. (CH3)2NH / H2SO4 (CH3)2NH+ / HSO4–
B. (CH3)2NH / (CH3)2NH+ H2SO4 / HSO4–
C. H2SO4 / HSO4– (CH3)2NH+ / (CH3)2NH
D. H2SO4 / (CH3)2NH (CH3)2NH+ / HSO4–
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BrønstedLowry
Amphoteric Substances
• Can act as either acid or base
– Can be either molecules or ions
e.g. Hydrogen carbonate ion:
– Acid
HCO3–(aq) + OH–(aq)  CO32–(aq) + H2O
– Base
HCO3–(aq) + H3O+(aq)  H2CO3(aq) + H2O
[Amphiprotic substances can donate or accept a
proton. This is a subtle but important difference from
the word amphoteric]
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Group
Problem
Which of the following can act as an
amphoteric substance?
A. CH3COOH
B. HCl
C. NO2–
D. HPO42–
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CHAPTER 16 Acids & Bases
Trends in
Acid/Base
Strength
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Acid/Base
Trends
Strengths of Acids & Bases
Strength of Acid
– Measure of its ability to transfer H+
– Strong acids
• React completely with water e.g. HCl and HNO3
– Weak acids
• Less than completely ionized e.g. CH3COOH and
CHOOH
Strength of Base classified in similar fashion:
– Strong bases
• React completely with water e.g. Oxide ion (O2–) and OH–
– Weak bases
• Undergo incomplete reactions
e.g. NH3 and NRH2 (NH2CH3, methylamine)
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Acid/Base
Trends
Strength in Water
• Strongest acid = hydronium ion, H3O+
– If more powerful H+ donor added to H2O
– Reacts with H2O to produce H3O+
Similarly,
• Strongest base is hydroxide ion (OH–)
– More powerful H+ acceptors
– React with H2O to produce OH–
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Acid/Base
Trends
Acid/Base Equilibrium
• Acetic acid (HC2H3O2) is weak acid
– Ionizes only slightly in water
HC2H3O2(aq) + H2O
H3O+(aq) + C2H3O2–(aq)
weaker acid
weaker base
stronger acid
stronger base
• Hydronium ion
– Better H+ donor than acetic acid
– Stronger acid
• Acetate ion
– Better H+ acceptor than water
– Stronger base
• Position of equilibrium favors weaker acid and base
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Group
Problem
In the reaction:
HCl + H2O → H3O+ + Cl–
which species is the weakest base ?
A. HCl
B. H2O
C. H3O+
D. Cl–
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Group
Problem
Group
Problem
Identify the preferred direction of the following
reactions:
H3O+(aq) + CO32–(aq)
HCO3–(aq) + H2O
Cl–(aq) + HCN(aq)
HCl(aq) + CN–(aq)
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Acid/Base
Trends
General Trends
• Stronger acids and bases tend to react with
each other to produce their weaker conjugates
– Stronger Brønsted acid has weaker
conjugate base
– Weaker Brønsted acid has stronger
conjugate base
• Can be applied to binary acids (acids made
from hydrogen and one other element)
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Acid/Base
Trends
Binary Acid Trends
Binary Acids = HnX
X = Cl, Br, P, As, S, Se, etc.
1.
Acid strength increases from left to right within
same period (across row)
– Acid strength increases as electronegativity
of X increases
e.g. HCl is stronger acid than H2S which is
stronger acid than PH3
– or
PH3 < H2S < HCl
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Acid/Base
Trends
Binary Acid Trends
Binary Acids = HnX
X = Cl, Br, P, As, S, Se, etc.
2. Acid strength increase from top to bottom
within group
– Acid strength increases as size of X and
bond length increases
e.g. HCl is weaker acid than HBr which is
weaker acid than HI
– or HCl < HBr < HI
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Group
Problem
Which is stronger?
• H2S or H2O
• CH4 or NH3
• HF or HI
Acid/Base
Trends
Oxoacid Trends
Oxoacids (HnX Om)
– Acids of H, O, and one other element
– HClO, HIO4, H2SO3, H2SO4, etc.
1. Acids with same number of oxygen atoms and differing X
a. Acid strength increases from
bottom to top within group
• HIO4 < HBrO4 < HClO4
b. Acid strength increases from left to
right within period as the electronegativity of the central
atom increases H3PO4 < H2SO4 < HClO4
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Acid/Base
Trends
Definition
Oxoacids (HnXOm)
2. For same X
– Acid strength increases with number of oxygen atoms
• H2SO3 < H2SO4
• More oxygens, remove more electron density from
central atom, weakening O—H bond make H more
acidic
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Group
Problem
Which is the stronger acid in each pair?
• H2SO4 or H3PO4
• HNO3 or H3PO3
• H2SO4 or H2SO3
• HNO3 or HNO2
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Group
Problem
Which corresponds to the correct order of
acidity from weakest to strongest acid ?
A. HBrO3, HBrO, HBrO2
B. HBrO, HBrO2, HBrO3
C. HBrO, HBrO3, HBrO2
D. HBrO3, HBrO2, HBrO
Acid/Base
Trends
Basicity
• Acid strength can be analyzed in terms of basicity
of anion formed during ionization
• Basicity
– Willingness of anion to accept H+ from H3O+
• Consider HClO3 and HClO4:
O
O
H
O
O
Cl
H
O
Cl
O
O
HClO3
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HClO4
30
Acid/Base
Trends
H
Basicity
O
O
O
Cl
O
H
O
Cl
O
O
HClO3
HClO4
• Lone oxygens carry most of the negative charge
– ClO4– has 4 O atoms, so each has –¼ charge
– ClO3– has 3 O atoms, so each has –1/3 charge
• ClO4– weaker base than ClO3–
– Thus conjugate acid, HClO4, is stronger acid
• HClO4 stronger acid as more fully ionized
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Group
Problem
Acid/Base
Trends
Organic Acid Trends
• Organic acid —COOH
• Presence of electronegative atoms (halide, nitrogen or
other oxygen) near —COOH group
– Withdraws electron density from O—H bond
– Makes organic acid, stronger acids
e.g.
CH3CO2H < CH2ClCO2H < CHCl2CO2H < CCl3CO2H
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Group
Problem
Which of the following is the strongest organic acid?
A
H
O
I
C
C
OH
B
H
O
Br
C
C
H
H
O
F
C
C
OH
E
H
O
Cl
C
C
H
H
O
H
C
C
H
H
D
H
OH
C
OH
OH
CHAPTER 16 Acids & Bases
Lewis Acid/Base
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Lewis
Acid/Base
Definition
• Broadest definition of species that can be
classified as either acid or base
• Definitions based on electron pairs
• Lewis acid
– Any ionic or molecular species that can accept
pair of electrons
– Formation of coordinate covalent bond
• Lewis base
– Any ionic or molecular species that can
donate pair of electrons
– Formation of coordinate covalent bond
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Lewis
Acid/Base
Lewis Neutralization
• Formation of coordinate covalent bond between electron
pair donor and electron pair acceptor
Addition Compound
• NH3BF3 = addition compound
– Made by joining two smaller molecules
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Lewis
Acid/Base
Lewis Acid-Base Reaction
Electrons in coordinate covalent bond come from O in
hydroxide ion
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Lewis
Acid/Base
Lewis Acids
1.
Molecules or ions with incomplete valence shells
e.g. BF3 or H+
2. Molecules or ions with complete valence shells, but with
multiple bonds that can be shifted to make room for
more electrons
e.g. CO2
3. Molecules or ions that have central atoms that can
expand their octets
– Capable of holding additional electrons
– Usually, atoms of elements in Period 3 and below
e.g. SO2
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Lewis
Acid/Base
Lewis Acid Example: SO2
O2–
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Lewis
Acid/Base
Lewis Bases
• Molecules or ions that have unshared electron pairs
and that have complete shells
– e.g. O2– or NH3
Lewis Definition is Most General
– All Brønsted acids and bases are Lewis acids and
bases
– All Arrhenius acids and bases are Brønsted acids
and bases
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Lewis
Acid/Base
Proton (H+) Transfer
H2O—H+ + NH3  H2O + H+—NH3
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Group
Problem
Identify the Lewis acid and base in the following:
• NH3 + H+
NH4+
• F– + BF3
BF4–
• SeO3 + O2–
SeO42–
Group
Problem
Which of the following species can act as
a Lewis base ?
A. Cl–
B. Fe2+
C. NO2–
D. O2–
CHAPTER 16 Acids & Bases
Acidity of Oxides
& Hydrates
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Acidic Metal
Ions
Acid-Base Properties of Elements &
their Oxides
Nonmetal oxides
– React with H2O to form acids
– Upper right hand corner of periodic table
– Acidic Anhydrides
– Neutralize bases
– Aqueous solutions red to litmus
– SO3(g) + H2O  H2SO4(aq)
– N2O5(g) + H2O  2HNO3(aq)
– CO2(g) + H2O  H2CO3(aq)
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Acidic Metal
Ions
Acid-Base Properties of Elements &
their Oxides
Metal oxides
– React with H2O to form hydroxide (Base)
– Group 1A and 2A metals (left hand side of periodic table)
– Basic Anydrides
– Neutralize acids
– Aqueous solutions blue to litmus
– Na2O(s) + H2O  2NaOH(aq)
– CaO(s) + H2O  Ca(OH)2(aq)
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Acidic Metal
Ions
Metal Oxides
• Solids at room temperature
• Many insoluble in H2O
• Why?
– Too tightly bound in crystal
– Can't remove H+ from H2O
– Do dissolve in solution of strong acid
– Now H+ free, can bind to O2– and remove
from crystal
Fe2O3(s) + 6H+(aq)  2Fe3+(aq) + 3H2O
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Group
Problem
What is the acid formed by P2O3
when it reacts with water ?
A. H2PO4
B. H2PO2
C. H3PO4
D. H3PO3
Acidic Metal
Ions
Metal Ions in Solution
• Exist with sphere of water molecules with their
negative poles directed toward Mn+
• Mn+(aq) + mH2O
Lewis Acid
Lewis Base
M(H2O)mn+(aq)
hydrated metal ion
= addition compound
– n = charge on metal ion
= 1, 2, or 3 depending on metal atom
– For now assume m = 1 (monohydrate)
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Acidic Metal
Ions
M(H2O)n+(aq) + H2O
Metal Hydrates are Weak Brønsted Acids
M(OH)n+(aq) + H3O+(aq)
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Acidic Metal
Ions
Metal Hydrates are Weak Brønsted Acids
• Electron deficiency of metal cations causes them to induce
electron density towards metal from water of hydration
• Higher charge density = more acidic metal
ionic charge
charge density 
ionic volume
• Acidity increases left to right across period
• Acidity decreases top to bottom down group
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Acidic Metal
Ions
Acidity of Hydrated Metal Ions
Degree to which M(H2O)mn+ produces acidic solutions
depends on:
1. Charge on Cation: As charge increases on Mn+, acidity
increases
– Increases metal ion’s ability to draw electron density to itself and
away from O—H bond
2. Cation’s Size: As size of cation decreases, acidity
increases
– Smaller, more concentrated charge
– Means greater pull of electron density from O—H bond
Net result: Very small, highly charged cations are very acidic
[Al(H2O)6]3+(aq) + H2O
[Al(H2O)5(OH)]2+(aq) + H3O+ (aq)
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Group
Problem
In the following list of pairs of ions, which
is the more acidic ?
Fe2+ or Fe3+; Cu2+ or Cu+; Co2+ or Co3+
A. Fe3+, Cu+, Co2+
B. Fe2+, Cu2+, Co3+
C. Fe3+, Cu2+, Co3+
D. Fe2+, Cu2+, Co2+
Acidic Metal
Ions
Trends in Acidity of Mn+
• Acidity increases up group (column) as cation size
decreases
• Acidity increases across period (row) as cation size
decreases
Alkali Metal Ions
(Li+, Na+, K+, Rb+, Cs+)
All weak
(+1, large size)
Be2+
Other Alkaline earth metals
(Ba2+, Ca2+ Sr2+, Mg2+)
Moderately weak
Very Weak
Transition metal ions, Al3(often +3,
Quite acidic
+4 charges)
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Group
Problem
Identify each of the following as acidic or basic and give their
reaction with water:
• P2O5
• MnO2
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