Download Section 5

Acids and Bases

Section 4
(Chapter 6, M,F&T)

Acid-Base and Donor-Acceptor
Chemistry
There are a variety of definitions for acids
and bases. Some definitions are quite
specific (and limited)
A few of the more popular acid-base
definitions are:
Arrhenius
Brønsted-Lowry
 Solvent system
 Lewis


Arrhenius Acids and Bases



Arrhenius acids are defined as substances which
increase the concentration of H3O+ ions when
added to water (e.g. H2SO4)
H2SO4 + H2O  HSO4- + H3O+
Arrhenius bases are substances that increase the
concentration of OH- ions when added to water
(e.g. NH3)
NH3 + H2O D NH4+ + OHIt is a definition that is limited to aqueous solutions
Brønsted Acids and Bases

A more general definition of acids and bases
that is defined as follows:
Brønsted acids are proton (H+) donors
 Brønsted bases are proton acceptors


The definition applies to all Arrhenius
cases, and even in non-aqueous solutions
HCl + H2O  H3O+ + ClNH3 + H2O D NH4+ + OHNH3 + HCl  NH4+ + Cl-
Structure and Brønsted Acidity


The ability of a Brønsted acid to donate a proton
will depend on the polarity of the H-X bond (in
most Brønsted acids, X = N, O, or a halogen)
Electron-withdrawing groups attached to X will
increase the quantity of partial positive charge on
the H-atom, making it more susceptible to
nucleophilic attack by a solvent (inductive effect)
O
O
H3C
O
acetic acid
H
F3C
O
H
Acids and Bases


There are a variety of definitions for acids
and bases. Some definitions are quite
specific (and limited)
A few of the more popular acid-base
definitions are:
Arrhenius
Brønsted-Lowry
 Solvent system
 Lewis


trifluoroacetic acid
O-H bond which is broken to yield H+ ion
1
Pauling’s Rules for Oxyacids
Structure and Brønsted Acidity

For oxyacids, acid strength increases with the
number of oxygen atoms:


more O-atoms, greater inductive effect
the stability of the conjugate base may also be the
driving force behind dissociation (resonance structures)


-
O
O
O Cl O H
O
O Cl O
+
+
H
+
H

-
O
O
O Cl O H
O Cl O
-
O Cl O H
+
O Cl O
H
+
+
Cl
Ka = big
O
O H
Cl
O
+
H
+
Ka = big

Ka = 1.1 x 10-2
Ka = 3.0 x 10-8
Solvent System Definition


The solvent system definition of acids and bases is
one that evolves from the autodissociation
reaction:
2 H2O D H3O+ + OHBy this definition, an acid is anything added to the
solvent that increases the concentration of the
cation of the autodissociation reaction (e.g. H3O+).
For example:
H2SO4 + H2O  H3O+ + HSO4-
Solvent System Definition


An acid is a substance that increases the concentration
of the cation of the autodissociation reaction

When SbF5 is added to a BrF3 solvent, the
following reaction occurs:
SbF5 + BrF3 D SbF6- + BrF2+
Thus SbF5 is an acid (in BrF3) by the solvent
system definition
A base is a substance that increases the concentration
of the anion of the autodissociation reaction


When KF is added to BrF3, the following reaction
takes place:
KF + BrF3 D K+ + BrF4Thus KF is a base by the solvent system definition
The solvent definition is also fairly general,
since many solvents are capable of
autodissociation:
2 NH3 D NH4+ + NH22 H2SO4 D H3SO4+ + HSO42 OPCl3 D OPCl2+ + OPCl42 BrF3 D BrF2+ + BrF4The last two equations don’t involve H+ ions
Solvent System Definition
Solvent System Definition

To predict the pKa of an oxyacid whose
formula can be written OpE(OH)q
pKa = 8 – 5p
hydrogen-free
oxygens
Where p is the number of hydrogen-free
oxygen atoms.
O
For polyprotic acids (for which q > 1), there
will be an increase in pKa of 5 units for
H
S
successive proton transfers
O
O
O
Sulfuric acid (O2S(OH)2, p = 2 and q = 2.
H
pKa1 ~ -2; pKa2 ~ +3
Even for protic solvents, this definition is more
useful than the Brønsted definition, since it treats
acidity not as an absolute property of the solute,
but must be specified in relation to the solvent
used.
 Example, for acetic acid (CH3COOH) in water:
CH3COOH + H2O D CH3COO- + H3O+
solute
 For acetic acid in H2SO4:
CH3COOH + H2SO4 D CH3COOH2+ + HSO4solute
 Thus acetic acid is an acid in water, but a base in
H2SO4

2
Lewis Acids and Bases

The definition proposed by Lewis is the
most general, and can be summarized by:



Lewis Acids and Bases

Lewis acids are electron-pair acceptors
Lewis bases are electron-pair donors

The following are examples:
H3O+ + OH-  2H2O
BF3 + Et2O  BF3OEt2
4NH3 + Cu2+  [Cu(NH3)4]2+
The Lewis acid-base reaction is driven by
the base’s ability to donate electrons to the
acid
Recognizing Lewis acids vs. Lewis bases is
not always easy, but
bases typically have lone pairs or negative
charges, while
 acids are often cations or may have empty
(acceptor) orbitals

Lewis Acids
Lewis Bases

Molecules possessing nitrogen atoms (amines,
imines, etc.) (e.g. ammonia, pyridine)
H


Cations (e.g. carbocations; electrophiles are thus
Lewis acids)
+
N
N
H
H

Molecules having oxygen atoms (e.g. water)
Includes metal ions (e.g. Fe3+)
3+
Fe
O
H

H

Anions (F-, C6H5COO-)
-
O
-
O
F
F


Molecules (or ions) that have
complete octets, but can
rearrange to accept more
electrons
Molecules that can handle
expanded octets (3rd period
elements and heavier) and can
accept additional e-’s
Closed-shell systems that can
accommodate more electrons
through p* orbitals
O

-
Ge
F
O
OH
O H
1. Adduct formation (base
donated e- pair to acid)
F
B
F
F
+
NC
2. Displacement reaction
F
F
F
N
CN
CN
CN substituents (cyano) are electron-withdrawing,
and lower the energy of the p* MO in this molecule
F
F
H
N H
H
+
H
B
F
F
F
F Ge F
F
2F
NC
2-
F
F
O
F
F
F
adduct formed with neutral
base indicated with arrow
-
C +
O -
B
Lewis Acid-Base Reaction Types
Lewis Acids

Molecules with empty (acceptor) orbitals (e.g. BF3)
and incomplete octets
F
F
H
H
H
B
N
F
H
H
H
B
+
N
N
F
H
H
N
CH3
3. Double displacement
H3C
Si
Br
CH3
CH3
+
AgCl
H3C
Si
Cl
CH3
+
AgBr
adduct formed with anionic
Lewis base indicated with line
3
Lewis Acids and Bases


The Acid-Base Interaction
The Lewis acid-base reaction is driven by
the base’s ability to donate electrons to the
acid
Recognizing Lewis acids vs. Lewis bases is
not always easy, but
bases typically have lone pairs or negative
charges, while
 acids are often cations or may have empty
(acceptor) orbitals

Electronic Factors
Hard Soft Acid-Base Concepts
Factors Influencing Acid-Base Reactions

There are four basic things which must be
considered in acid-base (donor-acceptor)
reactions:
1. The strength of the A-B bond (electronics)
2. The energy change involved in structural
rearrangements
3. Steric contributions
4. Solvent effects

Electron donors and acceptors tend to react in ways that
favor hard-hard and soft-soft interactions, proposed by
Pearson

Hard acids are small in size and/or highly charged (e.g. Li+,
Ti3+, BF3) (or whose d-electrons are relatively unavailable
for bonding) and bind preferentially to small/light basic
species
F- >> Cl- > Br- > IR2O >> R2S
R3N >> R3P

Soft acid species are polarizable, and are large, have low
charge if ionic (e.g. Ag+, BH3, Hg2+)
F- << Cl- < Br- < IR2O << R2S
R3N << R3P
Soft and Borderline Lewis Acids
- low or zero oxidation states, availability of d-electrons for p-bonding
4
Electronic Factors
Hard Soft Acid-Base Concepts
Hard-Hard

HSAB Guidelines
Soft-Soft
There is a greater separation
between the frontier orbitals
in a hard species than in a
soft species. Hard-hard
interactions have more ionic
character, while soft-soft
have more covalent
character.
Pearson’s Hardness Parameter
1
   I  A
2

Hard-hard and soft-soft interactions tend to
be favorable
Hard-hard creates strong interaction
because of ionic component
Hard-hard and soft-soft
interactions are favored
over hard-soft
Soft-soft interaction creates bonding MO
that is significantly
more stable (lower energy) than MO of
base (HOMO) or acid (LUMO)
Hardness
1
   I  A
2
Hard-Soft Acid Base Model
Pearson: favourable interactions:
Hard acid and hard base: ionic interactions dominant
Soft acid and soft base: covalent interactions dominant
Drago:
Quantitative treatment including parameters for electrostatic
and covalent contributions
A + B  AB
Hreaction (gas phase or in inert solvent)
-H = EAEB + CACB
5
Electronic Factors
Electronic Factors

HSAB Concepts
HSAB Concepts
Using HSAB guidelines, reactions between
acids and bases can be often be predicted
successfully (though not always)
Q: Why is AgI(s) very water-insoluble, but LiI very
water-soluble?
A: AgI is a soft acid-soft base combination, while LiI
is hard-soft. The interaction between Li+ and Iions is not strong.
Q: Is OH- or S2- more likely to form an insoluble
salt with a +3 transition metal ion?
A: The harder species will bind more strongly.
Between OH- or S2-, OH- is the harder species.
AgI(s) + H2O(l)  essentially no reaction
LiI(s) + H2O(l)  Li+(aq) + I-(aq)
Qualitative Analysis
Separation of Cations


In the separation of the group cations carried out
this year, HSAB rules are used to separate classes
of cations based on different hard and soft
interactions
soft and
Group II: Hg2+, Cd2+, Cu2+, Sn2+, Sb3+, Bi3+ borderline acids

Group III: Mn2+, Fe2+, Co2+, Ni2+, Zn2+, Al3+, Cr2+

H2S(g) D 2H+(aq) + S2-(aq)

Even at low S2- concentrations, the group II ions
precipitate (stronger interactions with the soft base, S2-)

Raising the pH increases the S2- concentration, which
allows the precipitation of group III ions

The group IV are then precipitated as hydroxides. These
cations are harder and prefer the hard base OH-.
borderline

Group IV: Ca2+, Mg2+, Ba2+, K+, NH4+
hard acids
Ambidentate Bases
GENERAL UNKNOWN

SCN- (thiocyanate) can interact through either its
S or N atom with Lewis acids. It can donate an
electron pair through more than one atom.

Interaction will be through the S-atom with a soft
acid, or through the N-atom when interacting with
hard acids.

Cr(III) interacts as Cr-NCS, while Pt(II) does so
as Pt-SCN
ACIDIC CONDITIONS
Decanted Solution
(Contains Group III & IV)
Precipitate containing
Group II Cations
BASIC CONDITIONS
Decanted Solution
Containing Group IV
Cations
Precipitate containing
Group III Cations
The soft and borderline cations are separated through
reaction with the soft base sulfide, S2-. Group II sulfides
are less soluble than group III, so in order to selectively
remove group II ions, a low pH is used:
6
Electronic Factors
Inductive Effects

Electron donating substituents
enhance base strength and
electron-withdrawing groups
enhance electron acceptor (acid)
strength
gas-phase
base strengths
Factors Influencing Acid-Base Reactions

P
Me
Me
Me
H
P
H
H
1. The strength of the A-B bond (electronics)
2. The energy change involved in structural
rearrangements
3. Steric contributions
4. Solvent effects
PMe3 stronger base than PH3
NMe3 > NHMe2 > NH2Me > NH3
strongest base
There are four basic things which must be
considered in acid-base (donor-acceptor)
reactions:
weakest base
This plays a role in bond lengths also
Me = methyl; alkyl, aryl groups are electron donating; F, CF3, CN, etc. are e- withdrawing
Structural Factors
Structural Rearrangement

Factors Influencing Acid-Base Reactions
In some cases, a center must adjust its hybridization in
order to accommodate the formation of a new bond
F
B
F
F
F
H
N H
+
F
B
N
F
H
opposite order to what is
expected for inductive effect
sp2

H
H
H
1. The strength of the A-B bond (electronics)
2. The energy change involved in structural
rearrangements
3. Steric contributions
4. Solvent effects
sp3

Order of Lewis acid strength for BX3 (X = halides) is
BF3 < BCl3 < BBr3

This is due to better p-orbital overlap in BF3 than in BCl3,
which is better than BBr3 (B-F bonds are shortest). Thus
more energy is needed to change from the sp2-hybridized
form of BF3.
Steric Factors
Influence of Solvent
Size/Bulk of Lewis Acid/Base

Solvent Properties
Bulky and/or large groups may interfere with
interaction between the donor and acceptor sites of
the base and acid
steric effect
N
H3C
CH 3
H3C
N
H3C
>
>
N
N
>
III
H 3C
>
H 3C
N
>
N
CH3
>
H3C
Since nearly all acid-base reactions occur in
solution, the properties of a solvent are critical to
the success or failure of a reaction.
There are five features of solvents that are
influential in acid-base reactions:

IV
reactions with H+ ions (inductive effect of alkyl donor enhances base strength;
bulkiness of t-butyl group in III offsets inductive effect
N


CH3
H 3C
II
I
There are four basic things which must be
considered in acid-base (donor-acceptor)
reactions:
CH3



N
H 3C
reactions with BF3 shows behavior that is influenced significantly by steric
effects of substituents

Usable temperature range Large temperature range desirable
Dielectric constant, e Important: ability to reduce attraction between ions
Solvent’s donor-acceptor properties Affect energies of reactants, products
Solvent’s protic acidity/basicity Will it protonate the reactants?
Nature and extent of autodissociation Solutes encounter not only
solvent molecules, but also
cations and anions of
autodissociation
7
Solvent Properties
Solvation Effects

Although in the gas phase, the amine bases exhibit the
following trend in base strength:
NMe3 > NHMe2 > NH2Me > NH3

In aqueous solution, the trend is
NHMe2 > NH2Me > NMe3 > NH3
and
NHEt2 > NH2Et ~ NEt3 > NH3


There are four basic things which must be
considered in acid-base (donor-acceptor)
reactions:
1. The strength of the A-B bond (electronics)
2. The energy change involved in structural
rearrangements
3. Steric contributions
4. Solvent effects
When the base reacts with water, the ammonium-type
conjugate acid produced is charged. The presence of three
methyl groups in NMe3 hampers the solvent’s ability to
solvate the charged ion (more H-atoms, more H-bonding),
making it less stable
Aquated Metal Ions
Factors Influencing Acid-Base Reactions

Me = methyl
Et = ethyl

The interaction of water
molecules with metal ions of high
charge and small size (or having a
high charge density) can lower the
pH of a solution, even though
there appears to be no proton
donor present
The base-acid interaction
weakens the O-H bond in
associated water molecules,
enabling H+ ions to be released
into solution
M
H
n+
+
O
H +
[M(H2O)6]n+(aq) + H2O(l) ⇌ H3O+(aq) + [M(H2O)5(OH)](n-1)+(aq)
Aquated Metal Ions

Smaller and highly charged cations (hard) like
Al3+, Fe3+, and Ti3+ are better at pulling away
electron density from water molecules than larger
ions, thus these aquated ions would be expected to
be quite acidic:
[Al(H2O)6]3+(aq) + H2O(l) ⇌ H3O+(aq) + [Al(H2O)5(OH)]2+(aq) pKa = 5.0
[Ti(H2O)6]3+(aq) + H2O(l) ⇌ H3O+(aq) + [Ti(H2O)5(OH)]2+(aq) pKa = 3.9
coordination complexes
For comparison, pKa for acetic acid is 4.74
8
Coordination Complexes
Aquated Ions: Interesting Cases

For [Cr(H2O)6]3+, formation of a dinuclear complex is
observed in basic solution (this also happens for Fe3+)
[Cr(H2O)6]3+(aq) + H2O ⇌ [Cr(H2O)5(OH)]2+ + H3O+(l)
2 Cr(H2O)5(OH)]2+(aq) ⇌ [(H2O)4Cr(mOH)2Cr(H2O)4]4+(aq) + 2H2O(l)
m denotes a “bridging” molecule.

When bases (“ligands”) interact with metal ions, a coordination
complex results. This interaction is created by a neutral or
anionic molecule (ligand) donating at least one electron pair to
the Lewis acid metal ion

The bonding in these complexes results from donation of an
electron pair by the ligand to the metal ion (coordinate covalent
bond)

Metal ions commonly coordinate four, six, or more ligands.

These types of complexes bridge the domains of inorganic
chemistry, organometallic chemistry (M-C bonds) and
biochemistry (porphyrins, proteins, etc.)
Bridging molecules (or bridging “ligands”) link Lewis acids
ferrocene
[Co(NH3)6]3+
Aquated Metal Ions
HEMOGLOBIN


HbO2 + CO ⇌ HbCO + O2
Lewis basicity:
O
O
C
K = 200
O
C
N
-
S2-
The interaction of water
molecules with metal ions of high
charge and small size (or having a
high charge density) can lower the
pH of a solution, even though
there appears to be no proton
donor present
The base-acid interaction
weakens the O-H bond in
associated water molecules,
enabling H+ ions to be released
into solution
M
H
n+
+
O
H +
[M(H2O)6]n+(aq) + H2O(l) ⇌ H3O+(aq) + [M(H2O)5(OH)](n-1)+(aq)
9
Coordination Complexes

When bases (“ligands”) interact with metal ions, a coordination
complex results. This interaction is created by a neutral or
anionic molecule (ligand) donating at least one electron pair to
the Lewis acid metal ion

The bonding in these complexes results from donation of an
electron pair by the ligand to the metal ion (coordinate covalent
bond)

Metal ions commonly coordinate four, six, or more ligands.

These types of complexes bridge the domains of inorganic
chemistry, organometallic chemistry (M-C bonds) and
biochemistry (porphyrins, proteins, etc.)
Geometrical Isomerism


Two species having the same molecular formula
and the same structural framework, but having
different spatial arrangements of atoms around a
central atom or double bond
Exists in




Square planar species: Pt(PPh3)2Cl2
Octahedral species: SnMe2F4, SH3F3
Trigonal bipyramidal species: Fe(CO)4PPh3
Double bonds (cis-, trans-): 2-butene
“LEWIS BASES”
mer- and fac- Isomerism
cis-, trans- Isomerism
X
X
cis-[Co(NH3)4Cl2]+
NH3
Co
NH3
M
Y
Cl
NH3
Y
Y
M
X
NH3
Y
Y
NH3
fac-MX3Y3
mer-MX3Y3
Cl
NH3
Co
Cl
NH3
trans-[Co(NH3)4Cl2]+
NH3
Cl
Cl
NH3
Cl
Pt
Pt
NH3
X
Y
X
Cl
X
NH3
cis-Pt(NH3)2Cl2
NH3
Cl
trans-Pt(NH3)2Cl2
10
Chelating Ligands


Some molecules/ions are capable of donating electron
pairs through more than one atom at once. This
interaction results in the formation of a chelate
(pronounced: key-late) ring
Chelating ligands tend to form very stable complexes
with metal ions. Some ligands are even capable of
forming more than one chelate ring (example EDTA:
ethylene diamine tetraacetic acid)
Nice to know: five- and six membered
rings tend to be the most stable, and
more chelate rings means more stable
How many chelate rings in this structure?
From Harris, Quantitative Chemical Analysis, 6 th Ed.
The Chelate Effect


Polydentate ligands form more stable complexes with
transition metal ions than monodentate ligands. They can
easily replace monodentate ligands in displacement reactions
For example, ethylene diamine (en) will replace ammonia in
[Cd(NH3)4]2+
[Cd(NH3)4]2+(aq) + 2en(aq) D [Cd(en)2]2+(aq) + 4NH3(aq)

The additional stability of a chelate complex over a
monodentate one is known as the chelate effect, and is
thermodynamic in origin
a bidentate ligand
H2N :
H2N
en =
: NH2
Chelate Effect

atoms in a ligand
Optical Isomerism
The chelate effect is a result of an entropy increase, and is
not so much an enthalpic effect:
G = H - TS
Cd2+(aq) + 4NH3(aq)  [Cd(NH3)4]2+(aq)
Ho = -52.5kJ/mol; So = -41.9 J/K.mol
Cd2+(aq) + 2en(aq)  [Cd(en)2]2+(aq)
Ho = -55.7 kJ/mol; So = +10.4 J/K.mol

NH2 denticity = # of donor
Mz+

Similar to carbon compounds, tetrahedral
complexes will also exhibit optical isomerism
(chiral complexes). Octahedral complexes
incorporating at least two bidentate ligands are
also chiral.
It is seen in the reaction below that four monodentate
ligands are displaced by two bidentate ligands, resulting in
a greater degree of disorder (So = +52.3 J/K.mol):
[Cd(NH3)4]2+(aq) + 2en(aq)  [Cd(en)2]2+(aq) + 4NH3(aq)
ENANTIOMERS
11
Optical Isomerism
Optical Isomerism

cis-complexes of this type exhibit this type of
isomerism, but not transmirror plane
N1
N1
Cl
N2
Cl
N2
Co
N3
Co
Cl
N3
Cl
N4
N4
Optical Activity
rotate 180o
A solution of one optical isomer will rotate plane-polarized light by +°
N1
Cl
N3
A solution of the other optical isomer will rotate it by -°
Co
N2
Cl
An equimolar mixture of the two isomers (racemic) will show no rotation
N4
What types of isomers can exist for
the following complexes?
“propeller complexes”
N
N
N
N
N
N

N

M
M
N

N
N
N
N

(no relation)





Lewis Acids and Bases


The Lewis acid-base reaction is driven by
the base’s ability to donate electrons to the
acid
Recognizing Lewis acids vs. Lewis bases is
not always easy, but
bases typically have lone pairs or negative
charges, while
 acids are often cations or may have empty
(acceptor) orbitals

[Ru(NH3)3(OH2)3]2+
Fe(CO)4Cl2
Ru(bpy)3
Ru(bpy)2Cl2
Ni(CO)2Br2
Cu(NH3)(OH2)BrCl
[Ru(tpy)2]2+
bpy
N
N
N
N
N
tpy
Polydentate Ligands

Other interesting polydentate ligands come
from the crown ether class of compounds
M+
12
Stability Constants of
Coordination Complexes
Crown Ethers

Consider the formation of ML6 (where L is a neutral
ligand) by the addition of L to an aqueous solution of
the cation:
[M(H2O)6]z+(aq) + 6L(aq) D [ML6]z+(aq) + 6H2O(l)
We can describe this formation reaction with a constant (like K):
 is the cumulative formation
constant (here, 6 ligands in one step)
We should break down the formation of this complex step-by-step,
since the coordination of each ligand involves
1. displacement of a water molecule
stepwise
2. coordination of the new ligand molecule
formation
For a metal cation of charge z+,
constants
[M(H2O)6]z+(aq) + L(aq)  [M(H2O)5L]z+(aq) + H2O(l)
[M(H2O)5L]z+(aq) + L(aq)  [M(H2O)4L2]z+(aq) + H2O(l)
[M(H2O)4L2]z+(aq) + L(aq)  [M(H2O)3L3]z+(aq) + H2O(l)
[M(H2O)3L3]z+(aq) + L(aq)  [M(H2O)2L4]z+(aq) + H2O(l)
[M(H2O)2L4]z+(aq) + L(aq)  [M(H2O)L5]z+(aq) + H2O(l)
[M(H2O) L5]z+(aq) + L(aq)  [ML6]z+(aq) + H2O(l)
K1
K2
K3
K4
K5
K6



6 
ML 
z
M ( H O ) L
6
2
z
6
6
We call K the stepwise stability (or formation)
constant. β is the cumulative stability (or formation)
constant
In contrast to solubility product constants and acid
dissociation constants, K is usually quite large
Thus, for
[M(H2O)6]n+(aq) + 6L(aq) D [ML6]n+(aq) + 6H2O(l)
β6 = K 1  K 2  K 3  K 4  K 5  K 6
or
log β6 = logK1 + logK2 + logK3 + logK4 + logK5 + logK6
where each K is calculated as
Kn =
éëM (H 2O)6-n Lnz+ ùû
éë M (H 2O)z+
ù
7-nû [ L ]
Stability Constants of F- Complexation
Stepwise stability constants for [Al(H2O)6-xFx](3-x)+ (x = 1 to 6)
A Possible Exam Question?
O
O
2-
C C
O
O
Consider the formation of a tris(oxalato)iron (III) salt
from [Fe(H2O)6]3+(aq). (oxalate = C2O42-)
Give expressions for the stepwise equilibria for the
formation of [Fe(ox)3]3- from Fe3+(aq) and ox2(log β1 = 7.54, log β2 = 14.59, log β3 = 20.00).
What are the numeric values of K1, K2, and K3?
Propose a reason for why K decreases in this series?
ox2-
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Answers



a) Fe3+(aq) = [Fe(H2O)6]3+(aq)
oxalate is a bidentate dianion (ox2-)
Stepwise formation of [Fe(ox)3]3-:
[Fe(H2O)6]3+(aq) + ox2-(aq) D [Fe(H2O)4(ox)]1+(aq) + 2H2O(l)
[Fe(H2O)4(ox)]1+(aq) + ox2-(aq) D [Fe(H2O)2(ox)2]1-(aq) + 2H2O(l)
[Fe(H2O)2(ox)2]1-(aq) + ox2-(aq) D [Fe(ox)3]3-(aq) + 2H2O(l)

b) β3 = K1K2K3, β2 = K1K2, and β1 = K1. So
K1 = 107.54 = 3.5 x 107
K2 = β2/K1 = 1.1 x 107
K3 = β3/K1K2 = 2.6 x 105

c) K will decrease as the charge of the
reactant complex decreases, since
electrostatic interaction will be less.
K1
K2
K3
The Hydrogen Bond – Donor-Acceptor Complex
Hydrogen Bonding in H2O
-
2d
d+
d+
H O
H O
Caused by: d+H
H
i) High POLARITY of the O-H bond
ii) Availability of unshared electrons on
oxygen
Limited to H and O?
NO! But need high electronegativities and
unshared electron pairs
H with N, O, F, (S, Cl)
Hydrogen Bonding
Do not confuse the phenomenon of
hydrogen bonding between molecules
with the bonds between O and H within
a molecule!
Hydrogen Bonding
14
The Hydrogen Bond

Definition of a ‘hydrogen bond’ is a moving target
A hydrogen bond is formed between an H atom attached to an
electronegative atom, and another electronegative atom that
possesses a lone pair of electrons.
The Hydrogen Bond
X−HB
Hydrogen bond formation has varying contributions from
three components:
An X−HB interaction is called a hydrogen bond if it
constitutes a local bond, and if X−H acts as a proton donor
towards Y.
1. An electrostatic component, from the polarity of the XH
bond.
The hydrogen bond is an attractive interaction between the
hydrogen from a group X−H and an atom or a group of atoms
B, in the same or different molecule(s), where there is evidence
of bond formation.
3. (London) dispersion forces.
2. A partial covalent character, and transfer of charge from
B to XH, from a donor-acceptor interaction.
Evidence for a Hydrogen Bond
XHB
XHB linear angle indicative of relatively strong H-bond,
short HB distance. Increased deviation from linearity, with
longer HB distances, indicates weaker H-bond.
Weakening, lengthening of XH bond, decreasing vibrational
frequency, formation of a new HB vibrational mode (IR,
Raman spectroscopies).
Deshielded H nucleus, strong downfield shift in 1H NMR
spectrum.
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Predicting H-Bond Strengths
Electrostatic Potential Map for Molecular Iodine I2
XHB ⇌ XHB ⇌ X¯HB+
- Competition between two acids, XH and HB+
pKa(XHB) = pKa(HX) - pKa(BH+)
Molecular Orbitals of I2
The Halogen Bond
Out-of-phase combination of
p-orbitals: * antibonding
LUMO
Near linear F-Cl-O due to alignment of acceptor * LUMO
In-phase combination of
p-orbitals: -bonding
Lengthening of F-Cl bond
16