section 4 · 2014. 3. 10. · 3 lewis acids and bases the definition proposed by lewis is the most...

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

-


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


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


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+


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




acid




charges, while


(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, Hg

2+)

F-

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




reactions:



rearrangements


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




reactions:



rearrangements


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




reactions:



rearrangements


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-MX3Y3 mer-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




acid




charges, while


(acceptor) orbitals

Polydentate Ligands

Other interesting polydentate ligands come

from the crown ether class of compounds

M+

http://en.wikipedia.org/wiki/File:18-crown-6-potassium-3D-vdW-A.png

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-nz+é

ëùû 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

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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

section 4 · 2014. 3. 10. · 3 lewis acids and bases the definition proposed by lewis is the most...

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