Chapter 15
Aqueous
Equilibria:
Acids and Bases
Properties of Acids and Bases
• Generally, an acid is a compound that releases hydrogen
ions, H+, into water.
– Blue litmus is used to test for acids. Blue litmus paper turns
red in the presence of hydrogen ions.
– Acids are generally sour in taste and can be corrosive.
• Generally, a base is a compound that releases hydroxide
ions, OH –, into water.
– Red litmus is used to test for bases. Red litmus paper turns
blue in the presence of hydroxide ions.
– Bases have a bitter taste and a slippery, soapy feel.
Chapter 15
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Definitions of an Acid and Base
• There are three definitions used to describe an
acid or a base:
Arrhenius Acids and Bases
Brønsted-Lowry Acids and Bases
Lewis Acids and Bases
Chapter 15
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Arrhenius Acids and Bases
• An Arrhenius acid is a substance that ionizes in water to
produce hydrogen ions (H+).
HA (aq) → H+ (aq) + A- (aq)
• An Arrhenius base is a substance that ionizes in water to
release hydroxide ions (OH-).
MOH (aq) → OH- (aq) + M+ (aq)
Chapter 15
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Brønsted-Lowry Acids & Bases
• The Brønsted-Lowry definitions of acids and bases are broader
than the Arrhenius definitions.
• A Brønsted-Lowry acid is a substance that donates a hydrogen ion
(H+) to any other substance.
– It is a proton donor.
– All Arrhenius acids are classified as acids by the B-L definition
Chapter 15
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Brønsted-Lowry Acids & Bases
• A Brønsted-Lowry base is a substance that accepts a
hydrogen ion (H+).
– It is a proton acceptor (So no OH- required!).
– All Arrhenius bases are classified as bases by the B-L definition
Chapter 15
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Conjugate Acid-Base Pairs
• When a pair of molecules are related by the loss or
gain of one H+, they are called a conjugate pair
• A conjugate acid-base pair is
the proton donor and the ion
formed by the loss of the
proton.
• A conjugate base-acid pair is
the proton acceptor and the ion
formed by the gain of the
proton.
Chapter 15
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Conjugate Acid-Base Pairs
HF (g) + H2O (l) ⇄ F- (aq) + H3O+ (aq)
NH3 (g) + H2O (l) ⇄ NH4+ (aq) + OH- (aq)
Chapter 15
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Conjugate Acid-Base Pairs
• Let’s practice writing and identifying conjugate pairs
What is the conjugate acid or base for the following:
Acid
HCl
H2O
HNO2
HSO4-
Conjugate Base
Base
HCO3H2O
NH3
CN-
Conjugate Acid
Identify the acid and base and their conjugates:
HCN + H2O ⇄ CN- + H3O+
NO2- + H2Se ⇄ HSe- + HNO2
Chapter 15
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Brønsted-Lowry Acids & Bases
•Write
balanced equations for the dissociation of
each of the following Brønsted–Lowry acids in
water.
(a) H2SO4
(b) HSO4–
(c) H3O+
•For
the above reactions, identify the conjugate
pairs and the acid and base in each pair.
Chapter 15
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Strengths of Acids and Bases
• B-L Acids and bases have
varying strengths.
• The strength of an acid is
measured by the amount of
H+ that is produced for each
mole of acid that dissolves
– Strong Acids fully ionize in
water
– ALL Products
• The strength of a base is
measured by the degree of
dissociation in solution.
– Strong Bases fully dissociate in
solution
– ALL Products
Chapter 15
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Strong Acids and Bases
Strong Acids
Sulfuric Acid
H2SO4
Hydroiodic Acid
HI
Hydrobromic Acid
HBr
Hydrochloric Acid
HCl
Perchloric Acid
HClO4
Nitric Acid
HNO3
• Only a few acids are considered
strong acids.
• Other acids have varying degrees
of weakness
• You need to know these!!
• Bases made from Groups IA and IIA metals and hydroxide are
strong bases:
LiOH, NaOH, KOH, Ca(OH)2, etc.
• Most other bases are weak bases
NH3, NH4OH, Al(OH)3, etc.
Chapter 15
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Strengths of Conjugate Acids and Bases
Chapter 15
13
Strengths of
Conjugate Acids
and Bases
Adcock, J Chem Ed (2001) 78, 1495-1496.
Chapter 15
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7
Strengths of Conjugate Acids and Bases
•If
you mix equal concentrations of reactants and
products, which of the following reactions proceed to the
right and which proceed to the left?
HF (aq) + NO3-(aq) ⇆ HNO3(aq) + F-(aq)
NH4+ (aq) + CO3-(aq) ⇆ HCO3 (aq) + NH3(aq)
Chapter 15
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Lewis Acids and Bases
• The Lewis definition of acids and bases is the broadest of
the three.
– A Lewis Acid is an electron pair acceptor. These are generally
cations and neutral molecules with vacant valence orbitals.
– A Lewis Base is an electron pair donor. These are generally
anions and neutral molecules with available pairs of electrons
(lone pairs!)
Chapter 15
The bond formed is a coordinate covalent bond!
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Lewis Acids and Bases
• Examples of Lewis acids include:
– Halides of Group 3A elements
– Oxides of Nonmetals (SO2, CO2 and SO3)
δ−
H
O
δ+
O
H
+
S
O
Lewis Acid
O
H
Lewis Base
O
O
H
S
O
O
H
O
S
O
O
H
O
Sulfuric Acid
• The S=O bond is polar with a δ+ on the less
electronegative S.
• This atom attracts the lone pair from the oxygen of
water
Chapter 15
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Lewis Acids and Bases
•
Identify the Lewis acid and Lewis base in each of the
following reactions:
SnCl4(s) + 2 Cl–(aq) ⇆ SnCl62–(aq)
Hg2+(aq) + 4 CN–(aq) ⇆ Hg(CN)42–(aq)
Co3+(aq) + 6 NH3(aq) ⇆ Co(NH3)63+(aq)
SO2(aq) + OH-(aq) ⇆ HSO3-(aq)
Chapter 15
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Summary of Acids & Bases
Arrhenius Acids/Bases
Brønsted-Lowry Acids/Bases
Lewis Acids/Bases
Chapter 15
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Hydrated Protons
• In aqueous solutions, the bare proton (H+) is too reactive to exist
alone, so it bonds to the oxygen atom of a solvent water
molecule to yield the hydronium ion (H3O+).
• As all of the acids we will be dealing with will be in aqueous
solution, the symbols H+ and H3O+ are interchangable
Chapter 15
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Autoionization of Water
• Water undergoes an autoionization reaction.
H2O(l) + H2O(l) ⇄ H3O+(aq) + OH-(aq)
OR
H2O(l) ⇄ H+(aq) + OH-(aq)
• Only about 1 in 5 million water molecules is
present as ions so water is a weak electrolyte.
• The concentration of hydrogen ions, [H+], in
pure water is about 1 × 10-7 mol/L at 25°C.
Chapter 15
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Autoionization of Water
• The molar ratio of H+ to OH- in the reaction is 1 to 1, so
if the [H+] = 1 × 10-7 mol/L at 25°C, then the [OH-] must
also be 1 × 10-7 mol/L at 25°C:
H2O (l) ⇄ H+ (aq) + OH- (aq)
[H+] • [OH-] = (1 × 10-7)(1 × 10-7) = 1.0 × 10-14
• This value is the ion product constant of water, Kw.
KW = [H+][OH-]
• We can use this value to calculate [H+] or [OH-]
Chapter 15
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[H+] and [OH-] Relationship
• The concentration of OH– ions in a certain household
ammonia cleaning solution is 0.00250 M. Calculate the
concentration of H+ ions.
• Calculate the concentration of OH– ions in a HCl
solution whose hydrogen ion concentration is 1.30 M.
Chapter 15
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The pH Scale
• pH is a measure of the acidity of a solution
Acidic solutions:
[H+] > 1.0 x 10–7 M,
pH < 7.00
Neutral solutions:
[H+] = 1.0 x 10–7 M,
pH = 7.00
Basic solutions:
[H+] < 1.0 x 10–7 M,
pH > 7.00
Chapter 15
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12
The pH Concept
• Mathematically:
pH = – log [H+]
OR
pH = – log [H3O+]
•
Nitric acid (HNO3) is used in the production of fertilizer, dyes,
drugs, and explosives. Calculate the pH of a HNO3 solution having a
hydrogen ion concentration of 0.76 M.
•
The pH of a certain orange juice is 3.33. Calculate the H+ ion
concentration.
Chapter 15
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The pOH Concept
• pOH is similar to pH except it is a measure of basicity
of a solution.
– The lower the value, the more basic the solution
pOH = - log [OH-]
• There is also a relationship between pH and pOH:
pH + pOH = 14
• So, if you know the [H+] you can determine the pH,
[OH-] and pOH
Chapter 15
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pH, pOH, [H+] and [OH-] Calculations
[H+]
[OH-]
pH
pOH
Acidic, Basic or Neutral?
6.15 x 10-4 M
2.61
5.28 x 10-8 M
3.45
pH = - log [H+]
pOH = - log [OH-]
[H+] = 10–pH
[OH-] = 10–pOH
[H+] • [OH-] = KW
pH + pOH = 14.00
Chapter 15
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Measuring pH
• The approximate pH of a solution can be determined
by using an acid-base indicator
• An indicator is a chemical substance that changes
color in a specific pH range.
Chapter 15
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14
Measuring pH
• Indicators let you determine the pH of your solution within
approximately ± 1 pH unit.
• More accurate methods, such as a pH meter, are also available
in most labs
Chapter 15
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pH of Strong Acids and Bases
• We know that strong acids and bases fully dissociate in solution
• Therefore, to determine pH, you need to do the following:
– Write and balance the dissociation reaction
– Determine the amount of H+ or H3O+ (or OH- if it’s a base) based on the
concentration of the acid (or base) and the mole ratio.
– Calculate the pH using the [H+] or [H3O+] (or [OH-] if it’s a base)
• Calculate the pH of the following solutions:
0.050 M HClO4
6.0 M KOH
0.125 M H2SO4
0.0357 M Ba(OH)2
Chapter 15
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15
Equilibria in Solutions of Weak Acids
• We know that strong acids fully dissociate in solution
• Solutions of weak acids contain both ionized and non-ionized
species in equilibrium
HA (aq) + H2O (l) ⇆ H3O+ (aq) + A- (aq)
[H 3O + ][A − ]
Ka =
[HA]
• Ka is the Acid Dissociation Constant
• This value can also be expressed as a pKa:
pKa = -log Ka
Chapter 15
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Equilibria in Solutions of Weak Acids
Note that the
larger the Ka
value, the
stronger the acid
WHY??
Chapter 15
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Equilibria in Solutions of Weak Acids
(a) Arrange the three acids in order of increasing value of Ka.
(b) Which acid, if any, is a strong acid?
(c) Which solution has the highest pH, and which has the
lowest?
Chapter 15
Determination of the
pH of a Weak Acid
Here, a good rule of thumb:
If Ka is significantly smaller
(> 1000 time) than [HA] the
assumption can be made.
33
STEP 1
Write the balanced equation
for the reaction
STEP 2
Under the balanced equation,
set up your ICE table
STEP 3
Write the Ka expression and substitute
the equilibrium concentrations
into the expression
STEP 4
Solve the Ka expression for X. If you have
made an assumption, make sure it is
valid. If it is not, you must use the
quadratic equation!
STEP 5
Go back to your ICE table and calculate
the Equilibrium concentrations using X
STEP 6
Use the equilibrium concentration of H+
+
or H3O to determine pH
Chapter 15
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Determination of the pH of a Weak Acid
• Calculate the pH of a 0.036 M nitrous acid
(HNO2) solution.
• What is the pH of a 0.122 M monoprotic acid
whose Ka is 5.7 x 10–4?
• The pH of a 0.060 M weak monoprotic acid is
3.44. Calculate the Ka of the acid.
Chapter 15
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Percent Dissociation in Solutions of Weak Acids
• The Percent Dissociation is a measure of the strength
of an acid.
[H + ]
% Dissociation =
× 100%
[HA]
• Stronger acids have higher percent dissociation.
• Percent dissociation of a weak acid decreases as its
concentration increases.
• Calculate the % dissociation of HF (Ka = 3.5 x 10-4) in
0.050 M HF
Chapter 15
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Equilibria in Solutions of Weak Bases
• We know that strong bases fully dissociate in solution
• Solutions of weak bases contain both ionized and non-ionized
species in equilibrium
B (aq) + H2O (l) ⇆ BH+ (aq) + OH- (aq)
[OH − ][BH + ]
Kb =
[B]
• Kb is the Base Dissociation Constant
• This value can also be expressed as a pKb:
pKb = -log Kb
Chapter 15
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Equilibria in Solutions of Weak Acids
Again, the larger the Kb value, the stronger the base
WHY??
Chapter 15
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Determination of the pH of a Weak Base
• Strychnine (C21H22N2O2), a deadly poison used to kill
rodents, is a weak base with a Kb = 1.8 x 10-6.
Calculate the pH of a saturated solution of strychnine
(16.0 mg/100 mL).
Chapter 15
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Relationship between Ka and Kb
• Product of Ka and Kb: multiplying out the expressions
for Ka and Kb equals Kw.
Ka × Kb = Kw
• What is the Ka of an ammonia solution (Kb = 1.8 x 10-5)?
• Calculate the pH of a 0.26 M methylamine solution.
Chapter 15
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Polyprotic Acids
• Polyprotic acids yield more than one hydrogen ion per
molecule.
• Protons are lost sequentially, one at a time.
• The conjugate base of first step is acid of second step.
• Ionization constants (Ka) decrease as protons are removed.
−
H2CO3(aq) + H2O(l) ⇆ H3O+(aq) + HCO3-(aq) K a1 =
[HCO 3 ][H 3O + ]
= 4.3 x 10-7
[H 2 CO 3 ]
HCO3-(aq) + H2O(l) ⇆ H3O+(aq) + CO32-(aq)
[CO 3 ][H 3O + ]
= 5.6 x 10-11
−
[HCO3 ]
2−
K a2 =
Chapter 15
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Polyprotic Acids
• Calculate the concentration of all species present
in a 0.10 M solution of oxalic acid (C2H2O4).
Determine the pH of the solution.
Chapter 15
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Molecular Structure and Acid Strength
• The strength of an acid depends on its tendency to ionize.
– The better it is at ionizing, the stronger the acid
• For general acids of the type H–X:
– The stronger the bond between H and X, the weaker the acid.
– The more polar the bond, the stronger the acid.
• For the hydrohalic acids, bond strength plays the key role giving:
Chapter 15
HF < HCl < HBr < HI
43
Molecular Structure and Acid Strength
• For binary acids in the same group, H–A bond strength decreases with
increasing size of A, so acidity increases.
• For binary acids in the same row, H–A polarity increases with increasing
electronegativity of A, so acidity increases.
Chapter 15
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Molecular Structure and Acid Strength
• For oxoacids bond polarity is more important. If we consider
the main element (Y):
Y
O
H + H2O
Y
O - + H3 O +
• If Y is an electronegative element, or in a high oxidation state,
the Y–O bond will be more covalent.
• This causes the O–H bond more polar and the acid stronger.
• Think of it the oxygen trying to hold onto electrons on both
sides
Chapter 15
45
Molecular Structure and Acid Strength
• For oxoacids with different central atoms that are
from the same group of the periodic table and that
have the same oxidation number, acid strength
increases with increasing electronegativity.
Chapter 15
46
23
Molecular Structure and Acid Strength
• For oxoacids having the same central atom but different numbers
of attached groups, acid strength increases with increasing
central atom oxidation number.
• The number of oxygen atoms increases the positive charge on the
chlorine which weakens the O–H bond and increases its polarity.
Chapter 15
47
Molecular Structure and Acid Strength
•Predict
the relative strengths of the following groups of
acids:
(a) H2S and H2Se
(b) HClO, HBrO, and HIO.
(c) HNO3 and HNO2.
(d) H3PO3 and H3PO4.
Chapter 15
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Acid-Base Properties of Salt Solutions
• When a salt dissolves in solution, it dissociates into
cations and anions.
• Solutions of salts can be acidic, basic or neutral.
• Anions and cations from strong acids and bases do not
affect pH.
• Anions and cations from weak acids and bases do affect
the pH.
• So, to determine whether your salt solution will be acidic,
basic or neutral, you have to figure out whether the anion
and cation would come from a strong or weak acid/base.
Chapter 15
49
Salts that Form Neutral Solutions
• A solution of a salt containing the cation from a
strong base and the anion from a strong acid will
be neutral (pH = 7)
NaCl (s) → Na+ (aq) + Cl- (aq)
Cation of
Strong Base
(NaOH)
Anion of
Strong Acid
(HCl)
Both from strong
Acid/Base so this solution
will be neutral
Chapter 15
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Salts that Form Basic Solutions
• A solution of a salt containing the cation from a
strong base and the anion from a weak acid will
be basic (pH > 7)
KF (s) → K+ (aq) + F- (aq)
Cation of
Strong Base
(KOH)
Anion of
Weak Acid
(HF)
The anion is the conjugate of a
weak acid so it is a strong base,
making the solution basic
Chapter 15
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Salts that Form Acidic Solutions
• A solution of a salt containing the cation from a
weak base and the anion from a strong acid will
be acidic (pH < 7)
NH4Br (s) → NH4+ (aq) + Br- (aq)
Cation of
Weak Base
(NH4OH)
Anion of
Strong Acid
(HBr)
The cation is the conjugate of a
weak base so it is a strong acid,
making the solution acidic
Chapter 15
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Salts that Can Go Either Way
• For salt solutions containing the cation from a
weak base and the anion from a weak acid the pH
will be dependent on the Ka and Kb
NH4F (s) → NH4+ (aq) + F- (aq)
If Ka > Kb then pH is Acidic
If Ka < Kb then pH is Basic
If Ka ≈ Kb then pH is Neutral
Chapter 15
Cation of
Weak Base
(NH4OH)
Anion of
Weak Acid
(HF)
You have to compare the Ka of
NH4+ and the Kb of F- to
determine pH
53
Acid-Base Properties of Salt Solutions
Chapter 15
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27
Acid-Base Properties of Salt Solutions
• Another type of acidic cation is a hydrated cation of a
small, highly charged metal ion
• Metal Ion Hydrolysis:
Chapter 15
55
Acid-Base Properties of Salt Solutions
•
Predict whether the following solutions will be acidic, basic,
or nearly neutral:
(a) NH4I
(b) CaCl2
(c) KCN
(d) Fe(NO3)3 (e) NH4CN
• Calculate the pH of a 0.15 M solution of sodium acetate
(CH3COONa).
• Calculate the pH of a 0.24 M ammonium chloride solution
(ZnCl2). The Ka = 2.5 x 10-10 for Zn(H2O)62+.
Chapter 15
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