Unit 11: Acid-Base Equilibrium
Chapter 16 and 17
Problem Set
• Chapter 16: 17, 21, 37, 43, 45, 61,
65, 69, 77, 79, 101, 107
• Chapter 17: 19, 23, 27, 31, 41,
WarmUp
Write the Formula
1) nitric acid
2) hydrofluoric acid
3) hydrobromic acid
4) chloric acid
5) acetic acid
6) carbonic acid
7) tetraboric acid
(tetraborate = B4O72-)
• Name the Acid
1) HCl
2) H2S
3) H3PO4
4) HI
5) HNO2
6) H3P
7) H2SO4
WarmUp
Write the Formula
1) nitric acid
2) hydrofluoric acid
3) hydrobromic acid
4) chloric acid
5) acetic acid
6) carbonic acid
7) tetraboric acid
(tetraborate = B4O72-)
• Name the Acid
1) HCl
2) H2S
3) H3PO4
4) HI
5) HNO2
6) H3P
7) H2SO4
Arrhenius Definition
• Acids produce hydrogen ions in
aqueous solution.
• Bases produce hydroxide ions when
dissolved in water.
• Limits to aqueous solutions.
• Only one kind of base.
• NH3 ammonia could not be an
Arrhenius base.
Bronsted-Lowry Definitions
• And acid is an proton (H+) donor and
a base is a proton acceptor.
• Acids and bases always come in pairs.
• HCl is an acid..
• When it dissolves in water it gives its
proton to water.
• HCl(g) + H2O(l)
H3O+ + Cl• Water is a base makes hydronium ion
Neutralization Reactions
• Remember: Acid + Base → Water + Salt
• The salt’s name and formula will be
based off of the cation of the base and
the anion of the acid
• Example: HCl + NaOH ↔ H2O + NaCl
• Neutralization reactions are one way that
is used to produce pure salts.
• Amphoteric Substances: can act as either
an acid or a base
Titration
• Method used to perform
neutralization and to
determine concentrations of an
unknown acid or base.
• Titrant – added standard
solution
• End Point – indicator color
change
• Equivalence Point – [H+] =
[OH-]
Acid Base Calculations
• In neutral solutions, [H+] and [OH-] are
equal
• [H+] = [OH-] = 1 x 10-7 M
• For other solutions, when [H+] increases,
then [OH-] will decrease and vise versa
• [H+] x [OH-] = 1.0 x 10-14
• Kw = [H+] x [OH-] = 1.0 x 10-14
• This is the Ion-Product Constant for
Water
• Acidic Solutions: [H+] > [OH-]
• Basic Solutions: [H+] < [OH-]
Using pH Scale
•
•
•
•
•
•
pH: Power of Hydrogen
pH = 7 Neutral [H+] = 1 x 10-7
pH < 7 Acidic [H] > 1 x 10-7
pH > 7 Basic [H+] = 1 x 10-7
pH = -log [H+]
Example Calculate pH if [H+] = 1 x 10-7 M
pH = -log (1 x 10-7 M)
= -(log 1 + log 10-7)
= -(0 +(-7))
pH = 7 * The exponent shows you the pH!*
The pOH Scale
• pOH : Power of Hydroxide
• pOH = 7 Neutral[OH-] = 1 x 10-7
• pOH < 7 Basic [OH-] > 1 x 10-7
• pOH > 7 Acidic [OH-] < 1 x 10-7
• pOH = -log [OH-]
*Same process as before, only now you are
dealing with [OH-]
Other Key Equations
• pH + pOH = 14
• pH = 14 – pOH
• pOH = 14 – pH
Calculations: Complete
the following chart
pH
pOH
[H]
[OH]
2.3
4.5
3.4x10-5
4.6x10-9
7
AcidBase?
How do we know if something is
a strong or weak acid/base?
We must calculate a value for the ionic
dissociation for that substance.
ACID
Ka = acid dissociation constant: the ratio of
the concentration of the dissociated form of
the acid to the concentration of the
undissociated form.
BASE
Kb = base dissociation constant: the ratio of
the product of the conjugate acid and OH
concentrations to the concentration of the
conjugate base
Strong and Weak Acids and Bases
• Dependent on how they dissolve in water
• Strong Acids completely ionize in
aqueous solutions
• Weak acids only partially ionize in
aqueous solutions
• Strong bases dissociate completely into
metal ions and hydroxide ions in aqueous
solutions
• Weak bases react with water to form OH
and the conjugate acid of the base
**Memorize Strong Acids and Bases**
Memorize the 8 strong acids…
all others are weak
•
•
•
•
HCl hydrochloric acid
HNO3 nitric acid
HBr hydrobromic acid
HIO4 periodic acid
•
•
•
•
HI hydroiodic acid
H2SO4 sulfuric acid
HClO4 perchloric acid
HClO3 chloric acid
Writing the Ka Expression
• Example HCl + H2O ↔ H3O+ + Cl• First you need to write a Keq expression
• Second, water can be eliminated because
its concentration is a constant.
• Write the Ka expression for this acid:
CH3COOH + H2O ↔ H3O+ + CH3COO-
Significance of Ka
• Ka indicated the fraction of acid in the
ionized form
• Weak acids have small Ka values
• Strong acids have large Ka values
because their ionization is more
complete
Writing the Kb Expression
• Example NH3 + H2O ↔ NH4+ + OH-
• Remember:
• Kb tells us how weak bases compete for
OH- from strong bases
• The smaller the value for Kb, the weaker
the base
Relationships
•
•
•
•
KW = [H+][OH-]
-log KW = -log([H+][OH-])
-log KW = -log[H+]+ -log[OH-]
pKW = pH + pOH
• KW = 1.0 x10-14
• 14.00 = pH + pOH
• [H+],[OH-],pH and pOH
Given any one of these we can find the
other three.
Example
• Calculate the pH of 2.0 M acetic acid
HC2H3O2 with a Ka 1.8 x10-5
• Calculate pOH, [OH-], [H+]
A mixture of Weak Acids
•
•
•
•
The process is the same.
Determine the major species.
The stronger will predominate.
Bigger Ka if concentrations are
comparable
• Calculate the pH of a mixture 1.20 M HF
(Ka = 7.2 x 10-4) and 3.4 M HOC6H5
(Ka = 1.6 x 10-10)
Salt Hydrolysis
• Most salt solutions are neutral, but some
can be acidic or basic
• Salts like sodium chloride and potassium
sulfate are neutral
WHY?
• Neutral salts are made from strong acids
and strong bases
• Salt Hydrolysis – the cations or anions of
the dissolved salt remove H+ from, or
donate H+ to water
• A solution is either acidic or basic based
on H+ transfer.
Salt Cations/Anions Effect on pH
In water:
• If cation and anion do not react, then
pH = neutral
• If cation does not react, but anion does,
then hydroxide is produced,
pH = basic
• If cation reacts but anion does not, then
H are produced,
pH = acidic
If both react:
OH- and H3O+ both produced
1. Anions from strong acids will not
change pH (ex: Br-)
2. Anions from a weak acid will increase
pH (ex: CN-)
3. Cations from weak bases will
decrease pH (ex: CH3NH3+)
4. Group 1A ions will increase pH
5. Group 2A ions will decrease pH
What is the pH of a 0.140M
solution of sodium acetate?
C2H3O2- + HOH↔HC2H3O2 + OHI
C
E
0.140 M
-x
0.140 – x
0
0
+x
+x
x
x
But what is Kb? KbxKa = Kw
Kb = Kw/Ka = 1x10-14/1.8x10-5
Solve for x
x = 8.77x10-6 =[OH-]
pOH=5.06, pH = 8.94 so the solution is basic.
The Common Ion Effect
• When the salt with the anion of a
weak acid is added to that acid,
• It reverses the dissociation of the acid.
• Lowers the percent dissociation of the
acid.
• The same principle applies to salts
with the cation of a weak base..
• The calculations are the same as last
chapter.
Types of Acids
• Monoprotic Acied – one acidic hydrogen
• Polyprotic Acids- more than 1 acidic hydrogen
(diprotic, triprotic).
• Oxyacids - Proton is attached to the oxygen of
an ion.
• Organic acids contain the Carboxyl group COOH with the H attached to O
• Generally very weak.
• H2SO3 ↔ H+ + HSO3- Ka1 = 1.7 x 10-2
• HSO3- ↔ H+ + SO32Ka2 = 6.4 x 10-8
• Why is K smaller?
• It is always easier to remove the first proton
Structure and Acid-Base
Properties
• Any molecule with an H in it is a potential
acid.
• The stronger the X-H bond the less acidic
(compare bond dissociation energies).
• The more polar the X-H bond the stronger
the acid (use electronegativities).
• The more polar H-O-X bond -stronger
acid.
Strength of oxyacids
• The more oxygen hooked to the
central atom, the more acidic the
hydrogen.
• HClO4 > HClO3 > HClO2 > HClO
• Remember that the H is attached to
an oxygen atom.
• The oxygens are electronegative
• Pull electrons away from hydrogen
Strength of oxyacids
Electron Density
Cl
O
H
Strength of oxyacids
Electron Density
O
Cl
O
H
Strength of oxyacids
Electron Density
O
Cl
O
O
H
Strength of oxyacids
Electron Density
O
O
O
Cl
O
H
Calculate the Concentration
• Of all the ions in a solution of 1.00 M
Arsenic acid H3AsO4
Ka1 = 5.0 x 10-3
Ka2 = 8.0 x 10-8
Ka3 = 6.0 x 10-10
Sulfuric acid is special
• In first step it is a strong acid.
• Ka2 = 1.2 x 10-2
• Calculate the concentrations in a 2.0
M solution of H2SO4
• Calculate the concentrations in a 2.0 x
10-3 M solution of H2SO4
Hydrated metals
• Highly charged
metal ions pull the
electrons of
Al+3
surrounding water
molecules toward
them.
• Make it easier for
H+ to come off.
H
O
H
Acid-Base Properties of Oxides
• Non-metal oxides dissolved in water
can make acids.
• SO3 (g) + H2O(l)
H2SO4(aq)
• Ionic oxides dissolve in water to
produce bases.
• CaO(s) + H2O(l)
Ca(OH)2(aq)
Lewis Acids and Bases
• Most general definition.
• Acids are electron pair acceptors.
• Bases are electron pair donors.
F
H
B
F
F
:N
H
H
Lewis Acids and Bases
• Boron triflouride wants more
electrons.
F
H
B
F
F
:N
H
H
Lewis Acids and Bases
• Boron triflouride wants more
electrons.
• BF3 is Lewis base NH3 is a Lewis Acid.
H
F
F
F
B
N
H
H
Lewis Acids and Bases
( )
( )
+3
Al + 6
H
O
H
H
Al
O
H
6
+3
Buffers
• Solutions in which the pH remains
relatively constant when small amounts
of acid of base are added.
• Acidic Buffer – composed of a weak acid
and one of its salts
• Basic Buffer – composed of a weak base
and one of its salts
• Buffer Capacity – the amount of acid or
base that can be added to a buffer
solution before a significant change in
pH occurs
Buffers have the ability to acquire
or give a way a H+, which allows it
to maintain a standard pH
• Example: Ethanoic Acid/Ethanoate system
(Acetic Acid / Acetate)
CH3COO- + H+ ↔ CHCOOH
CH3COOH + OH- ↔ CHCOO- + H2O
Both reactions occur simultaneously to
maintain pH
pH of Buffers
• To determine pH of Buffers, we use
the Henderson-Hasselbalch Equation
• Calculate the pH of a solution that is
.50 M HAc and .25 M NaAc
(Ka = 1.8 x 10-5)
Calculate the pH of the
following mixtures
1. 0.75 M lactic acid
(HC3H5O3) and 0.25 M
sodium lactate (Ka = 1.4 x
10-4)
2. 0.25 M NH3 and 0.40 M
NH4Cl (Kb = 1.8 x 10-5)
Titration Curves
Strong acid with strong Base
Equivalence at pH 7
•
•
pH
7
mL of Base added
Weak acid with strong Base
 Equivalence at pH >7

pH
>7
mL of Base added
Strong base with strong acid
 Equivalence at pH 7

pH
7
mL of Base added
Weak base with strong acid
 Equivalence at pH <7

pH
<7
mL of Base added
Summary
• Strong acid and base just stoichiometry.
• Determine Ka, use for 0 mL base
• Weak acid before equivalence point
–Stoichiometry first
–Then Henderson-Hasselbach
• Weak acid at equivalence point Kb
• Weak base after equivalence - leftover
strong base.
Summary
• Determine Ka, use for 0 mL acid.
• Weak base before equivalence point.
–Stoichiometry first
–Then Henderson-Hasselbach
• Weak base at equivalence point Ka.
• Weak base after equivalence - leftover
strong acid.
Indicators
• Weak acids that change color when they
become bases.
• weak acid written HIn
• Weak base
• HIn
H+ + Inclear
red
• Equilibrium is controlled by pH
• End point - when the indicator changes
color.
Indicators
• Since it is an equilibrium the color change is
gradual.
• It is noticeable when the ratio of
[In-]/[HI] or [HI]/[In-] is 1/10
• Since the Indicator is a weak acid, it has a Ka.
• pH the indicator changes at is.
• pH=pKa +log([In-]/[HI]) = pKa +log(1/10)
• pH=pKa - 1 on the way up
Indicators
• pH=pKa + log([HI]/[In-])
= pKa + log(10)
• pH=pKa+1 on the way down
• Choose the indicator with a pKa one
less than the pH at equivalence point
if you are titrating with base.
• Choose the indicator with a pKa one
greater than the pH at equivalence
point if you are titrating with acid.
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Acid-Base Equilibrium

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