Acids and Bases
Chapter 15
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
Properties of:
Acids
Bases
1.
2.
3.
4.
5.
Sour taste
Feels like water
Conducts electricity
Turns litmus paper red
Reacts with bases to
produce salts and water
6. Reacts with some metals
and releases hydrogen gas
o
Can you think of a reaction that
this occurs?
6. React with carbonates to
form carbon dioxide
1.
2.
3.
4.
5.
Bitter taste
Feels slippery
Conducts electricity
Turns litmus paper blue
Reacts with acids to
produce salts and water
6. Not Reactive with
metals
REVIEW ACID NAMING!!!
Binary Acids
• Contains only two different elements
– Hydrogen and an electronegative element
(usually a halogen)
• Nomenclature:
hydro - _________ - ic acid
Binary Nomenclature
Oxyacid
• Contains hydrogen, oxygen, and a third element
(hydrogen with a polyatomic ion)
• Nomenclature:
Polyprotic Acids
• Monoprotic:
– Donates 1 proton
– Example:
• Triprotic:
– Donates 3 protons
– Example:
• Diprotic:
– Donates 2 protons
– Example:
H2SO4 (l ) + H2O(l )  H3O (aq ) + HSO4– (aq )

2–


HSO4– (aq ) + H2O(l ) 
H
O
(
aq
)
+
SO
 3
4 (l )
STRONG ACIDS AND BASES
AND DEFINITIONS
Strength of Acids and Bases
MEMORIZE THESE!
•
•
•
•
•
•
•
Strong Acids
HCl
HBr
HI
HNO3
H2SO4
HClO4
HClO3
•
•
•
•
•
•
•
•
Strong Bases
LiOH
NaOH
KOH
RbOH
CsOH
Ca(OH)2
Sr(OH)2
Ba(OH)2
Strong Acids and Bases are
strong electrolytes
H3O+ (aq) + NO3- (aq)
HNO3 (aq) + H2O (l)
KOH (s)
H 2O
K+ (aq) + OH- (aq)
Weak Acids and Bases are
weak electrolytes
HF (aq) + H2O (l)
NO2- (aq) + H2O (l)
H3O+ (aq) + F- (aq)
OH- (aq) + HNO2 (aq)
Arrhenius Definition
• Arrhenius Acid:
– Substance that produces H+ (H3O+) in water
• Arrhenius Base:
– Substance that produces OH- in water
Acid example
But, NH3 is a base and does not have (OH-)!
Bronsted-Lowry Definition
• Bronsted-Lowry Acid:
– Proton donor
• Bronsted-Lowry Base:
– Proton acceptor
Forms conjugate acid-base pairs
acid is a proton donor
base is a proton acceptor
base
acid
conjugate
acid
conjugate
base
Conjugate Acid-Base Pairs
and Strengths
•
H3O+ is the strongest acid that can exist in aqueous
solution.
•
The OH- ion is the strongest base that can exist in aqueous
solution.
•
The stronger the acid, the weaker the conjugate base.
•
The stronger the base, the weaker the conjugate acid.
Lewis Acid and Base
• Lewis Acid:
– Substance that can accept a pair of electrons
• Lewis Base:
– Substance that can donate a pair of electrons
••
•• + OH
••
acid base
H+
H+ +
acid
••
H
N H
H
base
••
H O H
••
H
+
H N H
H
Lewis Acids and Bases
+
F B
••
H
F
N H
F
H
acid
base
F
F B
F
H
N H
H
No protons donated or accepted!
15.12
Acid and Base Definitions
ACID STRENGTH
Acid Strength
• Acid Strength depends on how easily
you can produce H+ ions
H X
H+
+
X
Binary Acid Strength DOWN a Group
H X
The bigger X is,
the weaker the HX
bond is
H+ + X-
Then the stronger the acid!
HF << HCl < HBr < HI
Binary Acid Strength ACROSS a Period
H X
H+ + X-
The greater
electronegativity of X, the
more polar the HX bond, the
easier to break the HX bond
Then the stronger the acid!
H3P < H2S << HCl
Oxyacid Strength
Z
dO
d+
H
Z
O- + H+
Depends on:
•
Electronegaitivity of “Z”
•
Number of oxygens (oxidation state) on “Z”
1. Electronegativity
Oxyacids having different central atoms (Z)
and that have the same number of oxygens.
Acid strength increases with increasing electronegativity of Z
••
••
••
O
••
••
••
H O Cl O
•• •• • •
••
••
H O Br O
•• •• • •
••
••
••
••
O
Cl is more electronegative than Br
HClO3 > HBrO3
2. Oxidation Numbers
Oxyacids having the same central atom (Z)
but different number of oxygen atoms.
Acid strength increases with the number oxygen atoms around Z.
HClO4 > HClO3 > HClO2 > HClO
The more oxygens
= stronger acid
Polyprotic Acids
• The strength of a polyprotic acid and its
anions decreases with increasing
negative charge.
• Increasing negative charge makes it
more difficult to release H+.
Increase # of H’s = stronger acid
H3PO4 › H2PO4- › HPO42-
TYPES OF SALTS
Neutral Salts
Neutral Solutions:
Salts containing an alkali metal or alkaline earth metal
ion (except Be2+) and the conjugate base of a strong
acid (e.g. Cl-, Br-, and NO3-).
NaCl (s)
H 2O
Na+ (aq) + Cl- (aq)
Comes from a
strong base
(NaOH)
Comes from a
strong acid
(HCl)
Basic Salts
Basic Solutions:
Salts derived from a strong base and a weak acid.
NaC2H3O2 (s)
H 2O
Na+ (aq) + C2H3O2- (aq)
Comes from a
strong base
(NaOH)
Comes from a
weak acid
(HC2H3O2)
Acidic Salts
Acid Solutions:
Salts derived from a strong acid and a weak base.
NH4Cl (s)
H 2O
NH4+ (aq) + Cl- (aq)
Comes from a
weak base
(NH3)
Comes from a
strong acid
(HCl)
Basic Oxides
• Metal oxides that react with water to form a base
Na2O + H2O
2 NaOH
Acidic Oxides
• Nonmetal oxides that react with water to form acids:
SO3 + H2O  H2SO4
Amphoteric Oxides
Intermediate element oxides (below stair step) that can
act as both an acid and a base
•React with acids to produce salt and water
Al2O3 (s) + 6 HCl (aq)
2 AlCl3(aq) + 3 H2O (l)
•React with bases to produce salt
Al2O3 (s) + 2 NaOH (aq) + 3 H2O (l)
2 Na Al(OH)4 (aq)
Oxides Summary
Basic Oxides
Amphoteric Oxides
Acidic Oxides
Metallic elements
Intermediate elements
Nonmetallic elements
Note: There are some exceptions but you will not be responsible for knowing these
PH CALCULATIONS
Water is Amphoteric
acts as both an acid and a base
H+ (aq) + OH- (aq)
H2O (l)
autoionization of water
H
O
H
+ H
[H
O
H
]
H
+ H
H
base
H2O + H2O
acid
O
+
conjugate
acid
H3O+ + OHconjugate
base
O
-
Ion Product Constant of Water
H2O (l)
H+ (aq) + OH- (aq)
[H+][OH-]
Kc =
[H2O]
[H2O] = constant
Kc= Kw = [H+][OH-]
ion-product constant (Kw):
product of the molar concentrations of H+ and OH- ions at a
particular temperature.
At 250C
Kw = [H+][OH-] = 1.0 x 10-14
What is the concentration of OH- ions in a HCl solution
whose hydrogen ion concentration is 1.3 M?
Given:
Kw = [H+][OH-] = 1.0 x 10-14
[H+] = 1.3 M
Solution:
-14
K
1
x
10
w
-15 M
=
=
7.7
x
10
[OH-] =
[H+]
1.3
Calculating pH
pH = -log [H+]
[H+] = [OH-]
At 250C
[H+] = 1 x 10-7
pH = 7
acidic
[H+] > [OH-]
[H+] > 1 x 10-7
pH < 7
basic
[H+] < [OH-]
[H+] < 1 x 10-7
pH > 7
Solution Is
neutral
pH
[H+]
15.3
Significant Figures
• There must be as many significant figures to the right
of the decimal as there are in the number whose
logarithm was found.
– Example: [H3O+] = 1 × 10−7
one significant figure
pH = 7.0
Calculating pOH
pOH = −log [OH–]
pH + pOH = 14.0
The pH of rainwater collected in a certain region of the
northeastern United States on a particular day was 4.82.
What is the H+ ion concentration of the rainwater?
pH = -log [H+]
[H+] = 10-pH = 10-4.82 = 1.5 x 10-5 M
The OH- ion concentration of a blood sample is 2.5 x 10-7 M.
What is the pH of the blood?
pOH = -log [OH-] = -log (2.5 x 10-7) = 6.60
pH + pOH = 14.00
pH = 14.00 – pOH = 14.00 – 6.60 = 7.40
The Circle of pH
-log [H3O+]
pH
[ H3O+]
10(-pH)
pH+ pOH = 14
[ H3O+] [ OH-]= 10-14
-log [OH-]
[ OH-]
pOH
10(-pOH)
What is the pH of a 2 x 10-3 M HNO3 solution?
HNO3 is a strong acid – 100% dissociation.
HNO3 (aq) + H2O (l)
Start 0.002 M
End
0.0 M
H3O+ (aq) + NO3- (aq)
0.0 M
0.0 M
0.002 M
0.002 M
pH = -log [H+] = -log [H3O+] = -log(0.002) = 2.7
What is the pH of a 1.8 x 10-2 M Ba(OH)2 solution?
Ba(OH)2 is a strong base – 100% dissociation.
Start
End
Ba(OH)2 (aq)
Ba2+ (aq) + 2OH- (aq)
0.018 M
0.0 M
0.0 M
0.0 M
0.018 M
0.036 M
pH = 14.00 – pOH = 14.00 + log(0.036) = 12.56
Weak Acids (HA) and
Acid Ionization Constants
HA (aq) + H2O (l)
HA (aq)
H3O+ (aq) + A- (aq)
H+ (aq) + A- (aq)
[H+][A-]
Ka =
[HA]
Ka = acid ionization constant
Ka
weak acid
strength
What is the pH of a 0.5 M HF solution (at 250C)?
HF (aq)
H+ (aq) + F- (aq)
[H+][F-]
= 7.1 x 10-4
Ka =
[HF]
HF (aq)
Initial (M)
Change (M)
Equilibrium (M)
H+ (aq) + F- (aq)
0.50
0.00
0.00
-x
+x
+x
0.50 - x
x
x
What is the pH of a 0.5 M HF solution (at 250C)?
Continued…
x2
= 7.1 x 10-4
Ka =
0.50 - x
Ka 
x2
= 7.1 x 10-4
0.50
Ka << 1
0.50 – x  0.50
x2 = 3.55 x 10-4
[H+] = [F-] = 0.019 M
[HF] = 0.50 – x = 0.48 M
pH = -log [H+] = 1.72
x = 0.019 M
When can I use the approximation?
Ka << 1 0.50 – x  0.50
When the percent ionization is less than or equal to 5%
x = 0.019
0.019 M
x 100% = 3.8%
0.50 M
To calculate percent ionization:
[H+]
x 100%
Percent ionization =
[HA]0
Less than 5%
Approximation ok.
[HA]0 = initial concentration
Does the approximation
always work?
What is the pH of a 0.050 M HF solution (at 250C)?
Ka 
x2
= 7.1 x 10-4
0.050
0.006 M
x 100% = 12%
0.050 M
x = 0.0060 M
More than 5%
Approximation not ok.
Must solve for x exactly using quadratic equation.
x2
= 7.1 x 10-4
Ka =
0.05 - x
ax2 + bx + c =0
x = 0.0056
HF (aq)
Initial (M)
Change (M)
Equilibrium (M)
x2 + 0.00071x – 3.55 x 10-5 = 0
-b ± b2 – 4ac
x=
2a
x = - 0.0063
H+ (aq) + F- (aq)
0.05
0.00
0.00
-x
+x
+x
0.05 - x
x
x
[H+] = x = 0.0056 M
pH = -log[H+] = 2.25
Solving weak acid ionization problems:
1. Identify the major species that can affect the pH.
•
In most cases, you can ignore the autoionization of
water.
•
Ignore [OH-] because it is determined by [H+].
2. Use ICE to express the equilibrium concentrations in terms
of single unknown x.
3. Write Ka in terms of equilibrium concentrations. Solve for x
by the approximation method. If approximation is not valid,
solve for x exactly.
4. Calculate concentrations of all species and/or pH of the
solution.
What is the pH of a 0.122 M monoprotic acid whose
Ka is 5.7 x 10-4?
HA (aq)
Initial (M)
Change (M)
Equilibrium (M)
0.122
0.00
0.00
-x
+x
+x
0.122 - x
x
x
x2
= 5.7 x 10-4
Ka =
0.122 - x
Ka 
H+ (aq) + A- (aq)
x2
= 5.7 x 10-4
0.122
0.0083 M
x 100% = 6.8%
0.122 M
Ka << 1
0.122 – x  0.122
x2 = 6.95 x 10-5
x = 0.0083 M
More than 5%
Approximation not ok.
x2
= 5.7 x 10-4
Ka =
0.122 - x
ax2 + bx + c =0
x = 0.0081
HA (aq)
Initial (M)
Change (M)
Equilibrium (M)
x2 + 0.00057x – 6.95 x 10-5 = 0
-b ± b2 – 4ac
x=
2a
x = - 0.0081
H+ (aq) + A- (aq)
0.122
0.00
0.00
-x
+x
+x
0.122 - x
x
x
[H+] = x = 0.0081 M
pH = -log[H+] = 2.09
Weak Bases and
Base Ionization Constants
NH3 (aq) + H2O (l)
NH4+ (aq) + OH- (aq)
[NH4+][OH-]
Kb =
[NH3]
Kb is the base ionization constant
Kb
weak base
strength
Solve weak base problems like weak acids
except solve for [OH-] instead of [H+].
Ionization Constants of
Conjugate Acid-Base Pairs
HA (aq)
A- (aq) + H2O (l)
H2O (l)
H+ (aq) + A- (aq)
OH- (aq) + HA (aq)
H+ (aq) + OH- (aq)
KaKb = Kw
Ka
Kb
Kw
Acid-Base Equilibria and
Solubility Equilibria
Chapter 16
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
Buffer Solutions
buffer solution: ability to resist changes in pH upon the
addition of small amounts of either acid or base.
Must have:
1. A weak acid or a weak base
AND
2. The salt of the weak acid or weak base
Buffer Example
Consider an equal molar mixture of CH3COOH and CH3COONa
Add strong acid
H+ (aq) + CH3COO- (aq)
Add strong base
OH- (aq) + CH3COOH (aq)
CH3COOH (aq)
CH3COO- (aq) + H2O (l)
Which of the following are buffer systems?
(a) KF/HF
(b) KBr/HBr
(c) Na2CO3/NaHCO3
(a) HF is a weak acid and F- is its conjugate base
buffer solution
(b) HBr is a strong acid
not a buffer solution
(c) CO32- is a weak base and HCO3- is it conjugate acid
buffer solution
What are some examples of
buffers?
16.3
TITRATIONS
Titrations
Titration:
when a solution of accurately known concentration is added
gradually added to another solution of unknown concentration
until the chemical reaction between the two solutions is
complete.
Indicator – substance that changes color at (or near) the
equivalence point
Slowly add base
to unknown acid
UNTIL
The indicator
changes color
(pink)
Titration Points
• equivalence point: point at which the two
solutions used in a titration are present in
chemically equivalent amounts
• end point: point in a titration at which an
indicator changes color
Common Indicators and
pH Ranges
Which indicator do I choose?
• pH less than 7
– Indicators that change color at pH lower than 7 are used
to determine the equivalence point of strong-acid/weakbase titrations.
– strong-acid/weak-base titration = acidic.
• pH at 7
– Indicators that undergo transition at about pH 7 are used
to determine the equivalence point of strong-acid/strong
base titrations.
– strong acids/strong bases = salt solution with a pH of 7.
The titration curve of a strong acid with a strong base.
16.5
Which indicator(s) would you use for a titration of HNO2
with KOH ?
Weak acid titrated with strong base.
At equivalence point, will have conjugate base of weak acid.
At equivalence point, pH > 7
Use cresol red or phenolphthalein
Strong Acid-Strong Base Titrations
NaOH (aq) + HCl (aq)
OH- (aq) + H+ (aq)
H2O (l) + NaCl (aq)
H2O (l)
0.10 M NaOH added to 25 mL of 0.10 M HCl
Weak Acid-Strong Base Titrations
CH3COOH (aq) + NaOH (aq)
CH3COOH (aq) + OH- (aq)
CH3COONa (aq) + H2O (l)
CH3COO- (aq) + H2O (l)
At equivalence point (pH > 7):
16.4
Strong Acid-Weak Base Titrations
HCl (aq) + NH3 (aq)
H+ (aq) + NH3 (aq)
NH4Cl (aq)
NH4+ (aq)
At equivalence point (pH < 7):
16.4
Strong Acid-Strong Base Calcs
• For monoprotic acids and bases:
MAVA = MBVB
Normality
• The value for normality is always larger
than that of molarity.
• Normality takes into consideration the
number of hydrogens/hydroxides a
substance will produce when
dissociated.
NAVA = NBVB
Normality Example
25.0 mL of 1.5 M sample of H3PO4 is
neutralized by 22.5 mL of Zn(OH)2.
What is the concentration of Zn(OH)2?
Antacids Video with Methyl Orange
KSP
Solubility Equilibrium
AgCl (s)
Ag+ (aq) + Cl- (aq)
Ksp = [Ag+][Cl-]
Ksp is the solubility product constant
MgF2 (s)
Ag2CO3 (s)
Ca3(PO4)2 (s)
Mg2+ (aq) + 2F- (aq)
2Ag+ (aq) + CO32- (aq)
3Ca2+ (aq) + 2PO43- (aq)
Ksp = [Mg2+][F-]2
Ksp = [Ag+]2[CO32-]
Ksp = [Ca2+]3[PO33-]2
16.6
Solubility Definitions
Solubility (g/L) is the number of grams of solute dissolved in
1 L of a saturated solution.
Molar solubility (mol/L):number of moles of solute dissolved
in 1 L of a saturated solution.
What is the solubility of silver chloride in g/L ?
Ksp = 1.6 x 10-10
AgCl (s)
Ag+ (aq) + Cl- (aq)
Initial (M)
0.00
0.00
Change (M)
Equilibrium (M)
+s
+s
s
s
Ksp = [Ag+][Cl-] = 1.6 x 10-10
Ksp = s2
s = Ksp = 1.3 x 10-5
[Ag+] = 1.3 x 10-5 M
[Cl-] = 1.3 x 10-5 M
-5 mol AgCl
1.3
x
10
143.35 g AgCl
Solubility of AgCl =
x
= 1.9 x 10-3 g/L
1 L soln
1 mol AgCl
16.6
Q vs. Ksp
Dissolution of an ionic solid in aqueous solution:
Q < Ksp
Unsaturated solution
Q = Ksp
Saturated solution
Q > Ksp
Supersaturated solution
No precipitate
Precipitate will form
If 2.00 mL of 0.200 M NaOH are added to 1.00 L of
0.100 M CaCl2, will a precipitate form?
The ions present in solution are Na+, OH-, Ca2+, Cl-.
Only possible precipitate is Ca(OH)2 (solubility rules).
Is Q > Ksp for Ca(OH)2?
From CaCl2:
0.100 M CaCl2 and 1.00 L
= 0.100 mol Ca2+ / 1.002 L
From NaOH:
0.200 M NaOH and 0.002 L
= 0.0004 mol OH- / 1.002 L
[Ca2+]0 = 0.100 M
[OH-]0 = 4.0 x 10-4 M
Q = [Ca2+]0[OH-]02 = 0.10 x (4.0 x 10-4)2 = 1.6 x 10-8
Ksp = [Ca2+][OH-]2 = 8.0 x 10-6
Q < Ksp No precipitate will form
common ion effect: shift in equilibrium caused by the addition
of a compound having an ion in common with the dissolved
substance.
The presence of a common ion suppresses
the ionization of a weak acid or a weak base.
Consider mixture of CH3COONa (strong electrolyte) and
CH3COOH (weak acid).
CH3COONa (s)
Na+ (aq) + CH3COO- (aq)
CH3COOH (aq)
H+ (aq) + CH3COO- (aq)
common
ion
The Common Ion Effect and Solubility
The presence of a common ion decreases
the solubility of the salt.
What is the molar solubility of AgBr in pure water?
AgBr (s)
Ag+ (aq) + Br- (aq)
Ksp = 7.7 x 10-13
s2 = Ksp
s = 8.8 x 10-7
Solubility with Common Ion
What is the molar solubility of AgBr in 0.0010 M NaBr?
NaBr (s)
Na+ (aq) + Br- (aq)
[Br-] = 0.0010 M
Ag+ (aq) + Br- (aq)
AgBr (s)
[Br-] = 0.0010 + s  0.0010
[Ag+] = s
Ksp = 0.0010 x s
s = 7.7 x 10-10
s = 8.8 x 10-7
Pure water
>
s = 7.7 x 10-10
Solution
without common ion > with common ion
Solubility
with and without common ion
AgBr in water vs. AgBr in NaBr Solution
s = 8.8 x 10-7
>
s = 7.7 x 10-10
without common ion > with common ion
pH and Solubility
•
•
•
The presence of a common ion decreases the solubility.
Insoluble bases dissolve in acidic solutions
Insoluble acids dissolve in basic solutions
remove
add
Mg(OH)2 (s)
Mg2+ (aq) + 2OH- (aq)
Ksp = [Mg2+][OH-]2 = 1.2 x 10-11
Ksp = (s)(2s)2 = 4s3
4s3 = 1.2 x 10-11
s = 1.4 x 10-4 M
[OH-] = 2s = 2.8 x 10-4 M
pOH = 3.55 pH = 10.45
If the normal pH is 10.45…
remove
add
Mg(OH)2 (s)
Mg2+ (aq) + 2OH- (aq)
If you lower the pH less than 10.45
Lower [OH-]
OH- (aq) + H+ (aq)
H2O (l)
Increase solubility of Mg(OH)2
If you raise the pH greater than 10.45
Raise [OH-]
Decrease solubility of Mg(OH)2