Transcript Aqueous Equilibria: Acids and Bases

Chapter 14
• Aqueous
Equilibria:
Acids and Bases
Slide 1
Acid–Base Concepts
01
Arrhenius Acid: A substance which dissociates to
form hydrogen ions (H+) in solution.
HA(aq)  H+(aq) + A–(aq)
Arrhenius Base: A substance that dissociates in, or
reacts with water to form hydroxide ions (OH–).
MOH(aq)  M+(aq) + OH–(aq)
Slide 2
Acid–Base Concepts
02
Brønsted–Lowry Acid: Substance that can donate H+
• Brønsted–Lowry Base: Substance that can accept H+
•
•
Chemical species whose formulas differ only by one
proton are said to be conjugate acid–base pairs.
Slide 3
Strong vs. Weak acids
03
Slide 4
Hydrated Protons and Hydronium
Ions
HA(aq)
H1+(aq) + A1-(aq)
Due to high reactivity of the hydrogen ion, it is
actually hydrated by one or more water molecules.
[H(H2O)n]1+
n=1
H3O1+
n=2
H5O21+
n=3
H7O31+
n=4
H9O41+
For our purposes, H1+ is equivalent to H3O1+.
Slide 5
Acid–Base Concepts
Slide 6
Lewis Acid–Base Concepts
Slide 7
Acid–Base Concepts
05
•
A Lewis Acid is an electron-pair acceptor. These
are generally cations and neutral molecules with
vacant valence orbitals, such as Al3+, Cu2+, H+, BF3.
•
A Lewis Base is an electron-pair donor. These are
generally anions and neutral molecules with
available pairs of electrons, such as H2O, NH3, O2–.
•
The bond formed is called a coordinate bond.
Slide 8
Acid–Base Concepts
06
-
+
Slide 9
Lewis Acids and Bases
Lewis Acid: An electron-pair acceptor.
Lewis Base: An electron-pair donor.
Slide 10
Lewis Acids and Bases
Lewis Acid: An electron-pair acceptor.
Lewis Base: An electron-pair donor.
Slide 11
Acid–Base Concepts
•
Write balanced equations for the dissociation of
each of the following Brønsted–Lowry acids.
(a) H2SO4
•
07
(b) HSO4–
(c) H3O+
Identify the Lewis acid and Lewis base in each of
the following reactions:
(a) SnCl4(s) + 2 Cl–(aq) æ SnCl62–(aq)
(b) Hg2+(aq) + 4 CN–(aq) æ Hg(CN)42–(aq)
(c) Co3+(aq) + 6 NH3(aq) æ Co(NH3)63+(aq)
Slide 12
Dissociation of Water
•
01
Water can act as an acid or as a base.
H2O(l) æ H+(aq) + OH–(aq)
•
Kc = [H+][OH–]
•
This is called the autoionization of water.
H2O(l) + H2O(l)
æ H3O+(aq) + OH–(aq)
Slide 13
Dissociation of Water
•
02
This equilibrium gives us the ion product constant
for water.
Kw = Kc = [H+][OH–] = 1.0 x 10–14
•
If we know either [H+] or [OH–] then we can
determine the other quantity.
Slide 14
Dissociation of Water
•
03
The concentration of OH– ions in a certain household
ammonia cleaning solution is 0.0025 M. Calculate the
concentration of H+ ions.
•
Calculate the concentration of OH– ions in a HCl
solution whose hydrogen ion concentration is 1.3 M.
Slide 15
pH – A Measure of Acidity
•
01
The pH of a solution is the negative logarithm of the
hydrogen ion concentration (in mol/L).
pH = –log [H+],
[H+] = 10-pH
pH + pOH = 14
Acidic solutions:
[H+] > 1.0 x 10–7 M,
pH < 7.00
Basic solutions:
[H+] < 1.0 x 10–7 M,
pH > 7.00
Neutral solutions:
[H+] = 1.0 x 10–7 M,
pH = 7.00
Slide 16
pH – A Measure of Acidity
02
•
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.
•
The OH– ion concentration of a blood sample is
2.5 x 10–7 M. What is the pH of the blood?
Slide 17
pH – A Measure of Acidity
04
Color of Tea: Polyphenols, Thearubigins
Color of Red Cabbage: Anthocyanin
Slide 18
pH – A Measure of Acidity
04
Slide 19
ACID
CONJ. BASE
HClO4
HI
HBr
HCl
H2SO4
HNO3
H3O+
HSO4–
ClO4–
I–
Br –
Cl –
HSO4 –
NO3 –
H2O
SO42–
ACID
Increasing Acid Strength
Increasing Acid Strength
Strength of Acids and Bases
HSO4–
HF
HNO2
HCOOH
NH4+
HCN
H2O
NH3
03
CONJ. BASE
SO42–
F–
NO2 –
HCOO –
NH3
CN –
OH –
NH2 –
Slide 20
Strength of Acids and Bases
•
04
Stronger acid + stronger base 
weaker acid + weaker base
•
Predict the direction of the following:
HNO2(aq) + CN–(aq) æ HCN(aq) + NO2–(aq)
HF(aq) + NH3(aq) æ F–(aq) + NH4+(aq)
Slide 21
Acid Ionization Constants
•
01
Acid Ionization Constant: the equilibrium
constant for the ionization of an acid.
HA(aq) + H2O(l) æ H3O+(aq) + A–(aq)
•
Or simply:
HA(aq) æ H+(aq) + A–(aq)


[H ][A ]
Ka 
[HA]
Slide 22
Conjugate Base Ionization Const
HA(aq) + OH−(aq)
A- + H2O(l)
[HA] [OH−]
Kb =
[A-]


−
[H
][A
]
[HA]
[OH
]
Ka  K b =

[HA]
[A ]
= Kw
Ka  Kb = Kw
Slide 23
Acid Ionization Constants
ACID
HF
HNO2
C9H8O4 (aspirin)
HCO2H (formic)
C6H8O6 (ascorbic)
C6H5CO2H (benzoic)
CH3CO2H (acetic)
HCN
C6H5OH (phenol)
Ka
7.1 x 10 –4
4.5 x 10 –4
3.0 x 10 –4
1.7 x 10 –4
8.0 x 10 –5
6.5 x 10 –5
1.8 x 10 –5
4.9 x 10 –10
1.3 x 10 –10
02
CONJ. BASE
Kb
F–
NO2 –
C9H7O4 –
HCO2 –
C6H7O6 –
C6H5CO2 –
CH3CO2 –
CN –
C6H5O –
1.4 x 10 –11
2.2 x 10 –11
3.3 x 10 –11
5.9 x 10 –11
1.3 x 10 –10
1.5 x 10 –10
5.6 x 10 –10
2.0 x 10 –5
7.7 x 10 –5
Slide 24
Strength of Acids and Bases
03
(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?
K=
(42/2) = 8
12/5= 0.2
Very Large
Slide 25
Acid Ionization Constants
Determine the pH of 0.50 M HA
solution at 25°C. Ka = 7.1 x 10–4
• Initial
05
Change Equilibrium Table:.
+
HA (aq) æ H (aq) + A (aq)
0.50
Initial (M):
Change (M):
–x
Equilib (M): 0.50 – x
0.00
+x
x
0.00
+x
x
Slide 26
What is the pH of a 0.50 M Citric acid solution (at 250C)?
+][A-]
[H
= 7.1 x 10-4
Ka =
HA (aq)
H+ (aq) + A- (aq)
[HA]
HA (aq)
Initial (M)
Change (M)
Equilibrium (M)
H+ (aq) + A- (aq)
0.50
0.00
0.00
-x
+x
+x
0.50 - x
x
x
100•Ka < Co ?
x2
-4
0.50 – x  0.50
= 7.1 x 10 100 x 7.1 x 10-4
Ka =
0.50 - x
= 0.071 < 0.5
x2
Ka 
x2 = 3.55 x 10-4 x = 0.019 M
= 7.1 x 10-4
0.50
[H+] = [A-] = 0.019 M
[HA] = 0.50 – x = 0.48 M
pH = -log [H+] = 1.72
Slide 27
Acid Ionization Constants
•
06
pH of a Weak Acid (Cont’d):
1.
Substitute equilibrium concentrations into equilibrium
expression.
2.
If 100•Ka < Co then (C0 – x) approximates to (C0).
3.
The equation can now be solved for x and pH.
4.
If 100•Ka is not significantly smaller than Co the
quadratic equation must be used to solve for x and pH.
Slide 28
Acid Ionization Constants
•
The Quadratic Equation:
•
The expression must first be rearranged to:
07
2
ax  bx  c  0
•
The values are substituted into the quadratic and
solved for a positive solution to x and pH.
2
 b  b  4ac
x
2a
Slide 29
Acid Ionization Constants
HA(aq)
•
09
H1+(aq) + A1-(aq)
Percent Dissociation: 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.
Slide 30
Percent dissociation of a weak acid
decreases as its concentration
increases
•
Concentration Dependence:
Slide 31
Weak Bases:
Base Ionization Constants
01
Base Ionization Constant:
The equilibrium constant for the ionization of a base.
• The ionization of weak bases is treated in the same
way as the ionization of weak acids.
•
B(aq) + H2O(l) æ BH+(aq) + OH–(aq)
•
Calculations follow the same procedure as used for
a weak acid but [OH–] is calculated, not [H+].
Slide 32
Base Ionization Constants
BASE
Kb
C2H5NH2 (ethylamine) 5.6 x 10 –4
CH3NH2 (methylamine) 4.4 x 10 –4
C8H10N4O2 (caffeine) 4.1 x 10 –4
NH3 (ammonia)
1.8 x 10 –5
C5H5N (pyridine)
1.7 x 10 –9
C6H5NH2 (aniline)
3.8 x 10 –10
NH2CONH2 (urea)
1.5 x 10 –14
CONJ. ACID
C2H5NH3+
CH3NH3+
C8H11N4O2+
NH4+
C5H6N+
C6H5NH3+
NH2CONH3+
02
Ka
1.8 x 10 –11
2.3 x 10 –11
2.4 x 10 –11
5.6 x 10 –10
5.9 x 10 –6
2.6 x 10 –5
0.67
Note that the positive charge sits on the nitrogen.
(caffeine)
Slide 33
Base Ionization Constants
•
03
Product of Ka and Kb: multiplying out the
expressions for Ka and Kb equals Kw.
Ka  Kb = Kw
Slide 34
pH of Basic Solutions
What is the pH of a 0.15 M solution of NH3?
NH4+(aq) + OH−(aq)
NH3(aq) + H2O(l)
[NH4+] [OH−]
Kb =
= 1.8  10−5
[NH3]
Initially
At Equilibrium
[NH3], M
[NH4+], M
[OH−], M
0.15
0
0
0.15 - x
x
x
Slide 35
pH of Basic Solutions
1.8  10−5 =
(x)2
(0.15 - x )
100 x Kb < C0 ?
1.8  10−3< 0.15
0.15 –x = 0.15
(1.8  10−5) (0.15) = x2
2.7  10−6 = x2
1.6  10−3 = x2
Slide 36
pH of Basic Solutions
Therefore,
X = [OH−] = 1.6  10−3 M
pOH = −log (1.6  10−3)
pOH = 2.80
pH = 14.00 − 2.80
pH = 11.20
Slide 37
Diprotic & Polyprotic Acids
01
H2SO4
H3PO4
•
Diprotic and polyprotic acids yield more than one hydrogen
ion per molecule.
•
One proton is lost at a time. Conjugate base of first step is
acid of second step.
•
Ionization constants decrease as protons are removed.
Slide 42
Diprotic & Polyprotic Acids
ACID
H2SO4
HSO4–
C2H2O4
C2HO4–
H2SO3
HSO3–
H2CO3
HCO3–
H2S
HS–
H3PO4
H2PO4–
HPO42–
Ka
Very Large
1.3 x 10 –2
6.5 x 10 –2
6.1 x 10 –5
1.3 x 10 –2
6.3 x 10 –8
4.2 x 10 –7
4.8 x 10 –11
9.5 x 10 –8
1 x 10 –19
7.5 x 10 –3
6.2 x 10 –8
4.8 x 10 –13
CONJ. BASE
HSO4 –
SO4 2–
C2HO4–
C2O42–
HSO3 –
SO3 2–
HCO3–
CO3 2–
HS–
S 2–
H2PO4–
HPO42–
PO43–
02
Kb
Very Small
7.7 x 10 –13
1.5 x 10 –13
1.6 x 10 –10
7.7 x 10 –13
1.6 x 10 –7
2.4 x 10 –8
2.1 x 10 –4
1.1 x 10 –7
1 x 10 –5
1.3 x 10 –12
1.6 x 10 –7
2.1 x 10 –2
Slide 43
Molecular Structure
and Acid Strength
01
•
The strength of an acid depends on its tendency to
ionize.
•
For general acids of the type H–X:
•
1.
The stronger the bond, the weaker the acid.
2.
The more polar the bond, the stronger the acid.
For the hydrohalic acids, bond strength plays the
key role giving: HF < HCl < HBr < HI
567 kJ/mol for HF
299 kJ/mol for HI
Slide 44
Molecular Structure
and Acid Strength
•
02
The electrostatic potential maps show all the hydrohalic
acids are polar. The variation in polarity is less
significant than the bond strength which decreases
from 567 kJ/mol for HF to 299 kJ/mol for HI.
Slide 45
(pm)
Slide 46
Molecular Structure
and Acid Strength
03
•
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.
Slide 47
Molecular Structure
and Acid Strength
•
04
For oxoacids bond polarity is more important. If we
consider the main element (Y):
Y–O–H
•
If Y is an electronegative element, the Y–O bond will
pull more electrons, the O–H bond will be more polar
and the acid will be stronger.
Slide 48
Molecular Structure
and Acid Strength
•
05
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.
Slide 49
Polar Covalent Bonds
02
Pauling Electronegativities
Detailed List of Electronegativity;
http://environmentalchemistry.com/yogi/periodic/electronegativity.html
Slide 50
Molecular Structure
and Acid Strength
•
07
Oxoacids of Chlorine:
Slide 51
Molecular Structure
and Acid Strength
•
08
Predict the relative strengths of the following
groups of oxoacids:
a) HClO, HBrO, and HIO.
b) HNO3 and HNO2.
c) H3PO3 and H3PO4.
Slide 52
Acid-Base Properties of Salts
Slide 53
Strong bases
Strong bases:
• The following metals make strong hydroxy base
•
•
•
Alkali metal cations of group 1A
Alkaline earth metal cations of group 2A
except for Be
Slide 54
Acid–Base Properties of Salts
01
•
Salts that produce neutral solutions are those
formed from strong acids and strong bases.
•
Salts that produce basic solutions are those formed
from weak acids and strong bases.
•
Salts that produce acidic solutions are those
formed from strong acids and weak bases.
Slide 55
Salts That Contain
Cation from a Weak Base and anion from a Weak Base
The pH of an ammonium carbonate solution, (NH4)2CO3,
depends on the relative acid strength of the cation and
the relative base strength of the anion.
Is it acidic or basic?
Slide 56
Acid-Base Properties of Salts
Salts That Contain Acidic Cations and Basic Anions
(NH4)2CO3:
NH41+(aq) + H2O(l)
H3O1+(aq) + NH3(aq)
Ka
CO32-(aq) + H2O(l)
HCO31-(aq) + OH1-(aq)
Kb
Three possibilities:
• Ka > Kb: The solution will contain an excess of
H3O1+ ions , Acidic solution, (pH < 7).
• Ka < Kb: The solution will contain an excess of
OH1- ions, Basic solutions, (pH > 7).
• Ka ≈ Kb: The solution will contain approximately
equal concentrations of H3O1+ and OH1- ions
(pH ≈ 7).
Slide 57
Salts That Contain
Cation from a Weak Acid and anion from a Weak Base
(NH4)2CO3:
NH41+(aq) + H2O(l)
H3O1+(aq) + NH3(aq)
Ka
CO32-(aq) + H2O(l)
HCO31-(aq) + OH1-(aq)
Kb
Ka for NH41+ =
Kb for CO32- =
Kw
Kb for NH3
1.0 x 10-14
=
Kw
Ka for HCO31-
1.8 x 10-5
= 5.6 x 10-10
1.0 x 10-14
=
5.6 x 10-11
= 1.8 x 10-4
Basic, Ka < Kb
Slide 58
Acid-Base Properties of Salts
Slide 59
Hydrated Cation of Al3+
Slide 60
Acid–Base Properties of Salts
•
03
Metal Ion Hydrolysis:
Slide 61
Acid–Base Properties of Salts
•
04
Calculate the pH of a 0.020 M Al(NO3)3 solution
Ka = 1.4 x 10-5.
•
Predict whether the following solutions will be
acidic, basic, or nearly neutral:
(a) NH4I
(b) CaCl2
(c) KCN
(d) Fe(NO3)3
Slide 62