16.5
Summary of MQ2 Results:
Mean =124 (71 %)
Hi = 175
Lo = 32
Your scores will be posted on WebCT
Using Ka to Calculate pH
• Percent ionization is another method to
assess acid strength.
• For the reaction
HA(aq) + H2O(l)
H3O+(aq) + A-(aq)
% ionization =
[H + ]eqm
× 100
[HA]0
• Percent ionization relates the equilibrium H+
concentration, [H+]eqm, to the initial HA
concentration, [HA]0.
Strong Acids and Bases
Strong Acids
Strong Bases
16.6 Weak Acids
Calculating Ka from pH
Using Ka to Calculate pH
Polyprotic Acids
16.7 Weak Bases
Types of Weak Bases
16.8 Relationship Between Ka and Kb
16.9 Acid-Base Properties of Salt Solutions
16.10
Acid-Base Behavior and Chemical Structure
Factors That Affect Acid Strength
Binary Acids
Oxyacids
Carboxylic Acids
16.11
Lewis Acids and Bases
Hydrolysis of Metal Ions
Using Ka to Calculate pH
• The higher percent ionization, the stronger
the acid.
• Percent ionization of a weak acid decreases
as the molarity of the solution increases.
• For acetic acid, 0.05 M solution is 2.0 %
ionized whereas a 0.15 M solution is 1.0 %
ionized.
Example 1: In 0.0100 M solution, acetic acid, abbrev HOAc,
is 4.2 % ionized. What is pH and Ka?
pH = 3.4
Ka = 1.8 x 10-5
Example 2: The pH of a 0.115 M solution of chloroacetic acid
is 1.92. What is Ka for this acid?
Ka = 1.4 x 10-3
Example 3: Calculate the concentrations of all species present
in a solution that is initially 0.10 M HOCl, Ka = 3.5 x 10-8
[H3O+] = [OCl-] = 5.9 x 10-5
[HOCl] = 9.9994 x 10-2
1
Notice we can summarize and compare, including use of pK
CH3COOH
ClCH2COOH
HOCl
HCN
HNO2
10-5
1.8 x
1.4 x 10-3
3.5 x 10-8
4.0 x 10-10
4.5 x 10-4
4.74
2.85
7.46
9.40
3.35
Likewise for Weak Bases:
NH3
(CH3)NH2
(CH3)2NH
(CH3)3N
C5H5N (pyridine)
1.8 x 10-5
5.0 x 10-4
7.4 x 10-4
7.4 x 10-5
1.5 x 10-9
4.74
3.30
3.13
4.13
8.82
Example 4: The pH of a household ammonia solution is 11.50.
What is the molarity of ammonia in the solution? Kb = 1.8 x 10-5
[NH3] = 0.57 M
An “indicator is just an organic dye with different colors for
the acid (HIn) and its conjugate base (In-). Consider
bromothymol blue
HIn + H2O
(yellow)
=
H3O+ +
In(blue)
An “indicator is just an organic dye with different colors for
the acid (HIn) and its conjugate base (In-). Consider
bromothymol blue
HIn + H2O
(yellow)
Ka =
or
=
H3O+ +
In(blue)
[ H 3O + ][[ In − ]
= 7.9 x10 −8
[ HIn]
[ HIn] [ H 3O + ]
=
[ In − ]
Ka
When this is >10 we see blue,
when <0.1 we see yellow
Polyprotic Acids
• Polyprotic acids have more than one ionizable
proton.
• The protons are removed in steps not all at once:
H2SO3(aq)
H+(aq) + HSO3-(aq) Ka1 = 1.7 x 10-2
HSO3-(aq)
H+(aq) + SO32-(aq)
Ka2 = 6.4 x 10-8
• It is always easier to remove the first proton in a
polyprotic acid than the second.
• Therefore, Ka1 > Ka2 > Ka3 etc.
• Most H+(aq) at equilibrium usually comes from the
first ionization (i.e. the Ka1 equilibrium).
2
Ascorbic acid, vitamin C, can be abbreviated as H2Asc .
It has K1a = 7.9 x 10-5 and K2a = 1.6 x 10-12.
What is the pH of a 0.10 M solution?
H2Asc + H20 = H3O+ + HAsc-
1st
K1a = 7.95 x 10-5 = x2/(0.100-x) ≈ x2/0.100
x = 0.0028
check assumption??
• Weak bases remove protons from
substances.
• There is an equilibrium between the base and
the resulting ions:
Weak base + H2O
• Example:
NH4+(aq) + OH-(aq)
NH3(aq) + H2O(l)
2nd
i
δ
HAsc- + H2O =
0.0028
-y
H3O+ +
0.0028
+y
Asc-2
0
+y
K2a = 1.6 x 10-12 = y(0.0028+y)y / (0.0028-y) ≈ y
y = [Asc-2] = 1.6 x 10-12
check [H3O+]
conjugate acid + OH-
• The base dissociation constant, Kb, is defined
as
[ NH +4 ][OH - ]
Kb =
[ NH 3 ]
Types of Weak Bases
• Bases generally have lone pairs or negative
charges in order to attack protons.
• Most neutral weak bases contain nitrogen.
• Amines are related to ammonia and have one
or more N-H bonds replaced with N-C bonds
(e.g., CH3NH2 is methylamine).
• Anions of weak acids are also weak bases.
Example: OCl- is the conjugate base of HOCl
(weak acid):
ClO-(aq) + H2O(l)
HClO(aq) + OH-(aq) Kb = 3.3 x 10-7
Relationship Between Ka and Kb
Relationship Between Ka and Kb
• We need to quantify the relationship between
strength of acid and conjugate base.
• When two reactions are added to give a third,
the equilibrium constant for the third reaction
is the product of the equilibrium constants for
the first two:
Reaction 1 + reaction 2 = reaction 3
has
K1 × K2 = K3.
• For a conjugate acid-base pair
Ka × Kb = Kw
• Therefore, the larger the Ka, the smaller the
Kb. That is, the stronger the acid, the weaker
the conjugate base.
• Taking negative logarithms:
pKa + pKb = pKw
3
Consider the following interesting reaction:
Consider the following interesting reaction:
HF
+
CN -
Æ
HCN
+
HF + CN-
F-
Which is the stronger acid? the stronger base?
Ka = 6.6 x 10-4
HF + H2O Æ H3O + + F and subtract
HCN + H2O Æ H3O + + CN - Ka‘ = 6.17 x 10-16
+
H3O + + F -
H3O + + CN - Æ HCN + H2O
HCN + F-
K = Ka(HF)/K’a(HCN) = 1.1 x 106
Two important acid equilibria:
HF + H2O Æ
=
Now consider CN- in water:
CN- + H2O = HCN + OH-
Ka = 6.6 x 10-4
K = 1 / Ka’ = 1.6 x 1015
giving our original reaction with K = Ka / Ka’ = 1.1 x
106
Kb(CN- ) = ?
Kb = Kw/Ka = (1.0 x 10-14) / 4.0 x 10-10)
= 2.5 x 10-5 = Kb(CN- )
!!!!
16.5
Summary of MQ2 Results:
Mean =122 (69 %)
Hi = 175 (4)
Lo = 21
Your scores are posted on WebCT
AcidAcid-Base Properties of Salt Solutions
• To determine whether a salt has acid-base
properties we use:
– Salts derived from a strong acid and strong base
are neutral (e.g. NaCl, Ca(NO3)2).
– Salts derived from a strong base and weak acid
are basic (e.g. NaOCl, Ba(C2H3O2)2).
– Salts derived from a weak base and strong acid
are acidic (e.g. NH4Cl, Al(NO3)3).
– Salts derived from a weak acid and weak base
can be either acidic or basic. Equilibrium rules
apply!
Strong Acids and Bases
Strong Acids
Strong Bases
16.6 Weak Acids
Calculating Ka from pH
Using Ka to Calculate pH
Polyprotic Acids
16.7 Weak Bases
Types of Weak Bases
16.8 Relationship Between Ka and Kb
16.9 Acid-Base Properties of Salt Solutions
16.10 Acid-Base Behavior and Chemical Structure
Factors That Affect Acid Strength
Binary Acids
Oxyacids
Carboxylic Acids
16.11 Lewis Acids and Bases
Hydrolysis of Metal Ions
AcidAcid-Base Properties of Salt Solutions
• Nearly all salts are strong electrolytes.
• Therefore, salts exist entirely of ions in solution.
• Acid-base properties of salts are a consequence of
the reaction of their ions in solution.
• The reaction in which ions produce H+ or OH- in
water is called hydrolysis.
• Anions from weak acids are basic.
• Anions from strong acids are neutral.
• Anions with ionizable protons (e.g. HSO4-) are
amphoteric.
4
Consider the effect of dissolving some NaOAc in water:
NaOAc(s) -----Æ Na+ + OAc-
Examples
but
OAc- + H3O+ = HOAc + H20
H20 + H2O
=
H3O+ + OH-
K = 1/Ka
Kw
so overall
OAc- + H2O = HOAc + OH-
Kb = usual expression
but we can rearrange to find that
Kb = Kw/Ka = (1.0 x 10-14) / (1.8 x 10-5) = 5.6 x 10-5
Calculate [OH-], pH, and % hydrolysis of a 0.10 M soln of NaCN
Contrast:
OAc- + H3O+ = HOAc + H2O
[HCN] = [OH-] = 1.6 x 10-3 M
[CN-] = (0.10 – 1.6 x 10-3) ≈ 0.10 M
1/Ka
pH = 14.00- pOH = 14.00 – 2.80 = 11.20
% hydrolysis = 1.6 %
OAc- + H2O = HOAc + OH-
Kb
Similar calculations for a solution of 0.10 M soln of NaOAc
would give [OH-] = 7.5 x 10-6
result of values of Ka for HOAc (1.8 x 10-5) and HCN (4.0 x 10-10)
Consider a 0.10 M solution of NaOAc,
Ka = 1.8 x 10-5 for HOAc
Consider a solution prepared by adding enough NaOCl to water
to make 2.00 L of solution with a pH = 10.50. How many moles
of NaOCl were added? Ka (HOCl) = 3.5 x 10-8
NaOCl Æ OCL - + H2O = HOCl + OH –
Kb = Kw / Ka = (3.5 x 10-8) / (1.0 x 10-14) = 2.86 x 10-7
[OCl- ]i = [NaOCl]i = 0.35 M
therefore 0.62 moles of NaOCl were added
5
AcidAcid-Base Behavior and Chemical Structure
Factors That Affect Acid Strength
Consider H-X. For this substance to be an acid
we need:
• H-X bond to be polar with Hδ+ and Xδ- (if X is
a metal then the bond polarity is Hδ-, Xδ+ and
the substance is a base),
• the H-X bond must be weak enough to be
broken,
• the conjugate base, X-, must be stable.
AcidAcid-Base Behavior and Chemical Structure
Binary Acids
• Acid strength increases across a period and
down a group.
• Conversely, base strength decreases across
a period and down a group.
• HF is a weak acid because the bond energy
is high.
• The electronegativity difference between C
and H is so small that the C-H bond is nonpolar and CH4 is neither an acid nor a base.
AcidAcid-Base Behavior and Chemical Structure
Oxyacids
• Oxyacids contain O-H bonds.
• All oxyacids have the general structure Y-O-H.
• The strength of the acid depends on Y and the
atoms attached to Y.
– If Y is a metal (low electronegativity), then the
substances are bases.
– If Y has intermediate electronegativity (e.g. I, EN
= 2.5), the electrons are between Y and O and
the substance is a weak oxyacid.
AcidAcid-Base Behavior and Chemical Structure
Strengths of H-O-Y Acids
(Oxyacids)
Oxyacids
– If Y has a large electronegativity (e.g. Cl, EN =
3.0), the electrons are located closer to Y than O
and the O-H bond is polarized to lose H+.
– The number of O atoms attached to Y increase
the O-H bond polarity and the strength of the acid
increases (e.g. HOCl is a weaker acid than
HClO2 which is weaker than HClO3 which is
weaker than HClO4 which is a strong acid).
Acid
EN of Y
Ka
pKa
HClO
3.2
3.0 x 10-8
7.52
HBrO
2.8
2.5 x 10-9
8.60
HIO
2.5
2.3 x 10-11
10.64
6
AcidAcid-Base Behavior and Chemical Structure
Carboxylic Acids
• These are organic acids which contain a
COOH group (R is some carbon containing
O
unit):
R
C
OH
AcidAcid-Base Behavior and Chemical Structure
Carboxylic Acids
• When the proton is removed, the negative
charge is delocalized over the carboxylate
anion:
O
R
C
O
O
R
C
O
• The acid strength increases as the number of
electronegative groups on R increases.
Lewis Acids and Bases
Lewis Acids and Bases
• Brønsted-Lowry acid is a proton donor.
• Focusing on electrons: a Brønsted-Lowry
acid can be considered as an electron pair
acceptor.
• Lewis acid: electron pair acceptor.
• Lewis base: electron pair donor.
• Note: Lewis acids and bases do not need to
contain protons.
• Therefore, the Lewis definition is the most
general definition of acids and bases.
• Lewis acids generally have an incomplete
octet (e.g. BF3).
• Transition metal ions are generally Lewis
acids.
• Lewis acids must have a vacant orbital (into
which the electron pairs can be donated).
• Compounds with π-bonds can act as Lewis
acids:
H2O(l) + CO2(g) → H2CO3(aq)
7
Lewis Acids and Bases
Hydrolysis of Metal Ions
• Metal ions are positively charged and attract water
molecules (via the lone pairs on O).
• The higher the charge, the smaller the metal ion
and the stronger the M-OH2 interaction.
• Hydrated metal ions act as acids:
Fe(H2O)63+(aq)
Fe(H2O)5(OH)2+(aq) + H+(aq) Ka = 2 x 10-3
• The pH increases as the size of the ion increases
(e.g. Ca2+ vs. Zn2+) and as the charge increases
(Na+ vs. Ca2+ and Zn2+ vs. Al3+).
8