Section 4 - StFXpeople.stfx.ca/dklapste/Chem245/c245_notes/Chem245_Acids and Ba… · 1 Section 4...
Transcript of Section 4 - StFXpeople.stfx.ca/dklapste/Chem245/c245_notes/Chem245_Acids and Ba… · 1 Section 4...
1
Section 4 (Chapter 6, M,F&T)
Acid-Base and Donor-Acceptor
Chemistry
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
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+ + OH-
It 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+ + Cl-
NH3 + H2O D NH4+ + OH-
NH3 + 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)
CH3
O
O
HF
3C O
O
H
acetic acid trifluoroacetic acidO-H bond which is broken to yield H+ ion
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
2
Cl
O
O
O
O H
Cl
O
OO H
Cl
O
O
O
O
Cl
O
OO
H+
H+
Cl OO H Cl OO
H+
Cl O H Cl O
H+
-
-
+
+
-+
-+
Ka = big
Ka = big
Ka = 1.1 x 10-2
Ka = 3.0 x 10-8
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)
Pauling’s Rules for Oxyacids
To predict the pKa of an oxyacid whose
formula can be written OpE(OH)q
pKa = 8 – 5p
Where p is the number of hydrogen-free
oxygen atoms.
For polyprotic acids (for which q > 1), there
will be an increase in pKa of 5 units for
successive proton transfers
Sulfuric acid (O2S(OH)2, p = 2 and q = 2.
pKa1 ~ -2; pKa2 ~ +3
OS
O
O
O
H
H
hydrogen-free
oxygens
Solvent System Definition
The solvent system definition of acids and bases is
one that evolves from the autodissociation
reaction:
2 H2O D H3O+ + OH-
By 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
The solvent definition is also fairly general,
since many solvents are capable of
autodissociation:
2 NH3 D NH4+ + NH2
-
2 H2SO4 D H3SO4+ + HSO4
-
2 OPCl3 D OPCl2+ + OPCl4
-
2 BrF3 D BrF2+ + BrF4
-
The last two equations don’t involve H+ ions
Solvent System Definition
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
When KF is added to BrF3, the following reaction takes place:
KF + BrF3 D K+ + BrF4-
Thus KF is a base by the solvent system definition
An acid is a substance that increases the concentration
of the cation of the autodissociation reaction
A base is a substance that increases the concentration
of the anion of the autodissociation reaction
Solvent System Definition
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+
For acetic acid in H2SO4:
CH3COOH + H2SO4 D CH3COOH2+ + HSO4
-
Thus acetic acid is an acid in water, but a base in
H2SO4
solute
solute
3
Lewis Acids and Bases
The definition proposed by Lewis is the
most general, and can be summarized by:
Lewis acids are electron-pair acceptors
Lewis bases are electron-pair donors
The following are examples:
H3O+ + OH- 2H2O
BF3 + Et2O BF3OEt2
4NH3 + Cu2+ [Cu(NH3)4]2+
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
Lewis Bases
Molecules possessing nitrogen atoms (amines,
imines, etc.) (e.g. ammonia, pyridine)
Molecules having oxygen atoms (e.g. water)
Anions (F-, C6H5COO-)
NHH
H
O
O
F
HO
H
N
-
-
Lewis Acids Cations (e.g. carbocations; electrophiles are thus
Lewis acids)
Includes metal ions (e.g. Fe3+)
Molecules with empty (acceptor) orbitals (e.g. BF3)
and incomplete octets
+
F BF
F
Fe3+
Lewis Acids
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
C
O
O
O H OH
O
O
+
-
-
-
-
GeFF
F
FF Ge
F
F
F
F F
F
2+
2-
NC CN
CNNC
CN substituents (cyano) are electron-withdrawing,
and lower the energy of the p* MO in this molecule
Lewis Acid-Base Reaction Types
1. Adduct formation (base
donated e- pair to acid)
2. Displacement reaction
3. Double displacement
B F
F
F
N
H
HH B
F
F
FN
H
H
H+
B
F
F
FN
H
H
H
N
B
F
F
FN N
H
H
H+
SiCH
3Br
CH3
CH3
AgCl SiCH
3Cl
CH3
CH3
AgBr+ +
adduct formed with neutral
base indicated with arrow
adduct formed with anionic
Lewis base indicated with line
4
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
The Acid-Base Interaction
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
Hard Soft Acid-Base Concepts
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- > I- R2O >> R2S R3N >> R3P
Soft acid species are polarizable, and are large, have low charge if ionic (e.g. Ag+, BH3, Hg2+)
F- << Cl- < Br- < I- R2O << R2S R3N << R3P
Electronic Factors
Soft and Borderline Lewis Acids
- low or zero oxidation states, availability of d-electrons for p-bonding
5
Hard Soft Acid-Base Concepts
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.
Electronic Factors
AI 2
1
Hard-Hard Soft-Soft
Pearson’s Hardness Parameter
HSAB Guidelines
Hard-hard and soft-soft interactions tend to
be favorable
Hard-hard creates strong interaction
because of ionic component
Soft-soft interaction creates bonding MO
that is significantly
more stable (lower energy) than MO of
base (HOMO) or acid (LUMO)
Hard-hard and soft-soft
interactions are favored
over hard-soft
AI 2
1
Hardness
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
6
HSAB Concepts
Using HSAB guidelines, reactions between
acids and bases can be often be predicted
successfully (though not always)
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.
Electronic Factors
HSAB Concepts
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 I-
ions is not strong.
Electronic Factors
AgI(s) + H2O(l) essentially no reaction
LiI(s) + H2O(l) Li+(aq) + I-(aq)
Qualitative Analysis
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
Group II: Hg2+, Cd2+, Cu2+, Sn2+, Sb3+, Bi3+
Group III: Mn2+, Fe2+, Co2+, Ni2+, Zn2+, Al3+, Cr2+
Group IV: Ca2+, Mg2+, Ba2+, K+, NH4+
soft and
borderline acids
borderline
hard acids
Separation of 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:
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-.
GENERAL UNKNOWN
Decanted Solution
(Contains Group III & IV)
Precipitate containing
Group II Cations
Decanted Solution
Containing Group IV
Cations
Precipitate containing
Group III Cations
ACIDIC CONDITIONS
BASIC CONDITIONS
Ambidentate Bases
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
7
Inductive Effects
Electron donating substituents
enhance base strength and
electron-withdrawing groups
enhance electron acceptor (acid)
strength
Electronic Factors
NMe3 > NHMe2 > NH2Me > NH3
strongest base weakest base
Me = methyl; alkyl, aryl groups are electron donating; F, CF3, CN, etc. are e- withdrawing
PMeMe
Me
PHH
H
PMe3 stronger base than PH3
This plays a role in bond lengths also
gas-phase
base strengths
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
Structural Rearrangement In some cases, a center must adjust its hybridization in
order to accommodate the formation of a new bond
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.
B F
F
F
N
H
HH B
F
F
FN
H
H
H+
Structural Factors
sp2 sp3
opposite order to what is
expected for inductive effect
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
Size/Bulk of Lewis Acid/Base
Bulky and/or large groups may interfere with
interaction between the donor and acceptor sites of
the base and acid
Steric Factors
NCH3
CH3
NCH3
N
CH3
CH3
CH3 N
NCH3
CH3
N NCH3
N
CH3
CH3
CH3
> > >
reactions with H+ ions (inductive effect of alkyl donor enhances base strength;bulkiness of t-butyl group in III offsets inductive effect
I II III IV
reactions with BF3 shows behavior that is influenced significantly by steric
effects of substituents
> > >
steric effect
Solvent Properties 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:
Usable temperature range
Dielectric constant, e
Solvent’s donor-acceptor properties
Solvent’s protic acidity/basicity
Nature and extent of autodissociation
Influence of Solvent
Large temperature range desirable
Important: ability to reduce attraction between ions
Will it protonate the reactants?
Affect energies of reactants, products
Solutes encounter not only
solvent molecules, but also
cations and anions of
autodissociation
8
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
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
Me = methyl
Et = ethyl
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
Aquated Metal Ions
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(H2O)6]n+(aq) + H2O(l) ⇌ H3O
+(aq) + [M(H2O)5(OH)](n-1)+(aq)
M O
H
Hn+
+
+
coordination complexes
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)
[Ti(H2O)6]3+
(aq) + H2O(l) ⇌ H3O+
(aq) + [Ti(H2O)5(OH)]2+(aq)
pKa = 5.0
pKa = 3.9
For comparison, pKa for acetic acid is 4.74
9
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(mOH)2Cr(H2O)4]
4+(aq) + 2H2O(l)
m denotes a “bridging” molecule. Bridging molecules (or bridging “ligands”) link Lewis acids
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.)
ferrocene
[Co(NH3)6]3+
HbO2 + CO ⇌ HbCO + O2 K = 200
Lewis basicity: O O C O C N-
S2-
HEMOGLOBIN Aquated Metal Ions
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(H2O)6]n+(aq) + H2O(l) ⇌ H3O
+(aq) + [M(H2O)5(OH)](n-1)+(aq)
M O
H
Hn+
+
+
10
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.)
“LEWIS BASES”
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
cis-, trans- Isomerism
Co
NH3
NH3
NH3
Cl
NH3
Cl
Co
NH3
NH3
Cl NH3
NH3
Cl
cis-[Co(NH3)4Cl2]+
trans-[Co(NH3)4Cl2]+
Pt
Cl NH3
Cl NH3
Pt
NH3 Cl
Cl NH3
cis-Pt(NH3)2Cl2 trans-Pt(NH3)2Cl2
mer- and fac- Isomerism
M XY
X
Y
X
YM XX
X
Y
Y
Y
fac-MX3Y3mer-MX3Y3
11
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)
From Harris, Quantitative Chemical Analysis, 6th Ed.
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?
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
NH2
NH2
NH2
NH2
:
:
Mz+
en =
a bidentate ligand
denticity = # of donor
atoms in a ligand
Chelate Effect
The chelate effect is a result of an entropy increase, and is not so much an enthalpic effect:
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
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)
G = H - TS
Optical Isomerism
Similar to carbon compounds, tetrahedral
complexes will also exhibit optical isomerism
(chiral complexes). Octahedral complexes
incorporating at least two bidentate ligands are
also chiral.
ENANTIOMERS
12
Optical Isomerism
cis-complexes of this type exhibit this type of
isomerism, but not trans-
Co
Cl
Cl
N1
N4
N2
N3
Co
Cl
Cl
N1
N4
N2
N3
Co
Cl
Cl
N1
N4
N2
N3
rotate 180o
mirror plane
Optical Isomerism
Optical Activity
A solution of one optical isomer will rotate plane-polarized light by +°
A solution of the other optical isomer will rotate it by -°
An equimolar mixture of the two isomers (racemic) will show no rotation
“propeller complexes”
M
N
N
N N
NN
M
N
N
NN
N N
(no relation)
What types of isomers can exist for
the following complexes?
[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+
N N
N
N
N
bpy
tpy
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
Polydentate Ligands
Other interesting polydentate ligands come
from the crown ether class of compounds
M+
13
Crown Ethers Stability Constants of
Coordination Complexes 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):
662
66
)( LOHM
MLz
z
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
2. coordination of the new ligand molecule
For a metal cation of charge z+,
[M(H2O)6]z+
(aq) + L(aq) [M(H2O)5L]z+(aq) + H2O(l) K1
[M(H2O)5L]z+(aq) + L(aq) [M(H2O)4L2]
z+(aq) + H2O(l) K2
[M(H2O)4L2]z+
(aq) + L(aq) [M(H2O)3L3]z+
(aq) + H2O(l) K3
[M(H2O)3L3]z+
(aq) + L(aq) [M(H2O)2L4]z+
(aq) + H2O(l) K4
[M(H2O)2L4]z+
(aq) + L(aq) [M(H2O)L5]z+
(aq) + H2O(l) K5
[M(H2O) L5]z+
(aq) + L(aq) [ML6]z+
(aq) + H2O(l) K6
where each K is calculated as
Kn =M (H2O)6-nLn
z+éë
ùû
M (H2O)7-n
z+éë
ùû L[ ]
stepwise
formation
constants
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 = K1 K2 K3 K4 K5 K6
or
log β6 = logK1 + logK2 + logK3 + logK4 + logK5 + logK6
Stability Constants of F- Complexation
Stepwise stability constants for [Al(H2O)6-xFx](3-x)+ (x = 1 to 6)
A Possible Exam Question?
Consider the formation of a tris(oxalato)iron (III) salt
from [Fe(H2O)6]3+(aq). (oxalate = C2O4
2-)
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?
C C
O
OO
O
2-
ox2-
14
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) K1
[Fe(H2O)4(ox)]1+(aq) + ox2-(aq) D [Fe(H2O)2(ox)2]1-(aq) + 2H2O(l) K2
[Fe(H2O)2(ox)2]1-(aq) + ox2-(aq) D [Fe(ox)3]
3-(aq) + 2H2O(l) K3
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.
The Hydrogen Bond – Donor-Acceptor Complex
H O
H
H O
H
2d-
d+d+
d+Caused by:
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 in H2O
Do not confuse the phenomenon of
hydrogen bonding between molecules
with the bonds between O and H within
a molecule!
Hydrogen Bonding Hydrogen Bonding
15
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.
An X−HB interaction is called a hydrogen bond if it
constitutes a local bond, and if X−H acts as a proton donor
towards Y.
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.
The Hydrogen Bond
Hydrogen bond formation has varying contributions from
three components:
1. An electrostatic component, from the polarity of the XH
bond.
2. A partial covalent character, and transfer of charge from
B to XH, from a donor-acceptor interaction.
3. (London) dispersion forces.
X−HB
Evidence for a Hydrogen Bond
XHB linear angle indicative of relatively strong H-bond,
short HB distance. Increased deviation from linearity, with
longer HB distances, indicates weaker H-bond.
Weakening, lengthening of XH bond, decreasing vibrational
frequency, formation of a new HB vibrational mode (IR,
Raman spectroscopies).
Deshielded H nucleus, strong downfield shift in 1H NMR
spectrum.
XHB
16
Predicting H-Bond Strengths
XHB ⇌ XHB ⇌ X¯HB+
pKa(XHB) = pKa(HX) - pKa(BH+)
- Competition between two acids, XH and HB+
Electrostatic Potential Map for Molecular Iodine I2
Molecular Orbitals of I2
In-phase combination of
p-orbitals: -bonding
Out-of-phase combination of
p-orbitals: * antibonding
LUMO
The Halogen Bond
Near linear F-Cl-O due to alignment of acceptor * LUMO
Lengthening of F-Cl bond