polar covalent bonds: electron distribution is...

Ch.2 Polar Bonds and Their Consequences

polar covalent bonds: electron distribution is unsymmetrical

Ionic Character

X Y X Y X Yδ+ δ- + -

symmetricalcovalent bond

polarcovalent bond ionic bond

2.1 Polar Covalent Bonds and Electronegativity


- EN difference: 0.3-2.0 (polar bond), >2.0 (ionic)

Elecrtonegativity (EN): the intrinsic ability of an atom to attract electrons in a covalent bond

H C N O F

PSi S Cl

Br

I

4.02.52.1 3.53.0

3.01.8 2.52.1

2.8

2.5

Li

Na

K

1.0

0.9

0.8

Be

Mg1.6

1.2Ca1.0

Cs0.7


δ-Cl

C

HH H

δ+ δ-

MgBr

C

HH H

δ+direction

of polarity

organometalliccompounds


inductive effects: shifting of electrons in σ-bond in response of the electronegativity of nearby atoms (atom's ability to polarize a bond); polarize bonds, chemical reactivity

Cl

C

HH H

electronegative atoms inductively

withdraw electrons

metals inductivelydonates electrons

MgBr

C

HH H


dipole moment (μ): net molecular polarity (vector sum); resulted from individual bond polarities and lone pair contributions; polar substances dissolve in polar solvents, nonpolar substances soluble in nonpolar solvents

µ = Q x r Q: charge at either end of dipoler: distance between the charges

1 debyes (D): 3.336 x 10-30 C . m

For one positive charge and one negative charge seperated by 100 pm

-B

A+100 pm

µ = Q x r = (1.60 x 10-19 C) x (100 x 10-12 m)(1D / 3.336 x 10-30 C.m)= 4.80 D

2.2 Polar Covalent Bonds and Dipole Moment


δ-Cl

C

HH H

δ+

µ = 1.87 D

C-Cl: 178 pm

µ = 4. 80 D x 1.78 = 8.54 D

For fully seperated chloromethane (if the C-Cl bond were ionic, C+ Cl-)

but the actual dipole moment is 1.87 D

x 100= 22%

therefore, the chlorine atom in chloromethane has an excess of about 0.2 electron,and the carbon has a deficiency of about 0.2 electron

1.87

8.54

δ-Cl

C

HH H

δ+

0.2 electron excess

0.2 electron deficient


nonbonding electron pair: stick out from positively charged nuclei

H

OH

netnetN

HH H

H

OH

N

HH H


NaCl 9.0

Dipole moment(D)

H3C NO

O3.46

CH3ClH2OCH3OH

H2C=N=N

1.871.851.701.50

NH3 9.0

CH4

CCl4CH3CH3

000

0BF3

0

Dipole moment(D)

Table 2.1 dipole moments of some compounds


symmetrical molecules: μ = 0 D


• related to bond polarity, dipole moment; common for atoms that have an apparently abnormal number of bonds

C NO

O

H

H

H

formal positive chargeformal negative charge

Nitromethane

2.3 Formal Charges


C C

H

HHH

4 valence electronsfor isolated carbon atom

N N HHH

5 valence electronsfor isolated nitrogen atom

electron counts- covalent bond: each atom owns one electron; single bond (1 e-); double bond (2 e-); triple bond (3 e-)

- non-bonding electrons: owned by the atom

8

2= 4 valence electrons

6

2+ 2 = 5 valence electrons


Formal Charge = # of valence electronsin free atom

# of valence electronsin bound atom

_

# of valenceelectrons

half ofbondingelectrons

# ofnonbonding

electrons

_ _=


C NO

O

H

H

H

Nitromethane

nitrogen valence electronsnitrogen bonding electronsnitrogen nonbonding electrons

= 5= 8= 0

oxygen valence electronsoxygen bonding electronsoxygen nonbonding electrons

= 6= 2= 6

For neutral molecules: sum of the formal charges equal to zero

Formal Charge = 5 - 8/2 - 0 = + 1

Formal Charge = 6 - 2/2 - 6 = - 1


C C C

+ 1 0 - 1

Table 2.2 A summary of formal charges on atoms

• dipolar molecule: neutral overall but have +/- charges on individual atoms; dipolar character of molecules → chemical reactivity

N N N

+ 1 0 - 1

O O O

+ 1 0 - 1

H3CS

CH3

O

Practice formal charge

Dimethyl sulfoxide

H3CS

CH3

O

H3CS

CH3

Oor

H2C N N

Diazomethane

H2C N N

H3C C N

Acetonitrile

H3C C N

H3C N C

Methyl isocyanide

H3C N C



• For some molecules simple Lewis structure cannot describe the actual structure correctly

C NO

O

H

H

H

Nitromethane

C NO

O

H

H

H

-

2.4 Resonance

For example, nitromethane has two equivalent N-O bonds, but one Lewis structure can't represent it

Both oxygen-nitrogen bonds are 122 pm in length, midway between the lengthOf a typical N-O single bond (130 pm) and a typical N=O double bond (116 pm)


resonance form: individual Lewis structures that represent one molecule


1. Individual resonance forms are imaginary, not real; The real structure is composite (or resonance hybrid) of different forms.

C NO

O

H

H

H

Nitromethane

C NO

O

H

H

H

-

2.5 Rules for Resonance Forms

2. Resonance forms differ only in the placement of their π or nonbonding electrons; Neither the position nor the hybridization of any atom changes from one resonance form to another.

A curved arrow shows movement of two electrons.


CC

CCC

CH

H

HH

H

H

CC

CCC

CH

H

HH

H

H

Benzene (two resonance forms)


H3C CH3

O strongbase

H3C C

O

H

H H3C C

O

H

H

3. Different resonance forms of a substance don't have to be equivalent; The real structure is composite (or resonance hybrid) of different forms.

For non-equivalent resonance forms, the actual structure of the resonance hybrid is close to the more stable form than the less stable form.

more stable form(negative charge on

electronegative oxygen)


4. Resonance forms must be valid Lewis structures and obey normal rules of valency; The octet rule still applies.

5. The resonance hybrid is more stable than any individual resonance form; The larger the number of resonance forms, the more stable a substance is.

C CO

O

H

H

H

Acetate ion (two resonance forms)

C CO

O

H

H

H

-C C

O

O

H

H

H

NOT a validresonance form

10 electrons on carbon


X

Y

Z

*

XY

Z* *XY

Z X

Y

Z

*

* = 0, 1, or 2 electronsmultiple bond (double or triple bond)

X

Y

Z

2.6 Drawing Resonance Forms

Any three atom grouping with a multiple bond has two resonance forms.


O O

H

O O

H

OO

H

OC

O

O_ _

Practice Resonance forms

OC

O

O_ _

OC

O

O_ _

OC

O

O_

_

OC

O

O_

_

Carbonate ion: CO32-


Practice Resonance forms

CC

CC

CH

HH

H H

H

H

Pentadienyl radical

CC

CC

CH

HH

H H

H

HC

CC

CC

HH

H

H H

H

HC

CC

CC

HH

H

H H

H

H

allylic radical

A half-headed curved arrow shows movement of one electron.



2.7 Acids and Bases: The BrØnsted-Lowry Definition

BrØnsted-Lowry Definition: proton donor/acceptorLewis Definition: electron pair donor/acceptor

BrØnsted-Lowry Acid: proton (H+) donorBrØnsted-Lowry Base: proton (H+) acceptor


+ H-B+ +

Acid Base Conjugateacid

Conjugatebase

H-A B A-

In general,

+ H2O +


Conjugatebase

CH3COOH -OH CH3COO-

+ NH3 +


Conjugatebase

H2O -NH2-OH


2.8 Acid and Base Strength

+ H3O+ +H-A H2O A-

Keq =[H3O+][A-]

[HA][H2O]

Ka = Keq [H2O] =[H3O+][A-]

[HA]

Acidity constant, Ka

in dilute solution [H2O] remains constant [H2O] = 55.6 M for pure water

pKa = -logKa

stronger acid: larger Ka, smaller pKaweaker acid: smaller Ka, larger pKa


Acid pKa Conjugate base

CH3CH2OHH2OHCNCH3COOHHFHNO3HCl

CH3CH2O-

HO-

CN-

CH3COO-

F-

NO3-

Cl-

16.0015.749.314.763.45- 1.3 - 7.0

strongeracid

strongerbase

Table 2.3 Relative strength of some common acids and their conjugate bases


- inverse relationship between the acid strength and its conjugate base strength

- a strong acid loses H+ easily; its conjugate base has little affinity for H+

- a weak acid loses H+ with difficulty; its conjugate base has high affinity for H+

for example, HCl is a strong acid; Cl- does not hold H+ tightly and is thus a weak base

+ H-B+ +


Conjugatebase

H-A B A-


2.9 Predicting Acid-Base Reactions from pKa Values

An acid with a lower pKa will go to the conjugate acid with a higher pKa.

+ H2O+

Acetic acid(pKa = 4.76)

CH3COOH -OH CH3COO-

Water(pKa = 15.74)

Hydroxide ion Acetate ion

strongeracid

strongerbase

weakeracid

weakerbase

Practice Acidity

C C HH + HO- C CH + HO-H

pKa = 15.74pKa = 25

?

+

Acetone(pKa ~ 19)

H3C CH3

ONa+NH2

- ?

C C HH + NH2- C CH + NH3

pKa = 35pKa = 25


+

Amonia(pKa ~ 35)

H3C CH2

ONH3

Na+


2.10 Organic Acids and Organic Bases• organic acids are two kinds: O-H and O=C-C-H

H3C OH

H3C O

Anion is stabilized by havingnegative charge on a highly electronegative atom

H3C O

OH

H3C O

O

H3C O

O

Anion is stabilized by having negative charge on a highly electronegative atom and by resonance

pKa = 15.54

pKa = 4.76


O H

H3C CH3

O O

H

pKa = 10 pKa = 9

H3C O

OH

pKa = 4.76

H3C CH2

OH

H3C CH2

O

H3C CH2

O

Anion is stabilized resonance and having negative charge on a highly electronegative atom

organic acids

pKa = 19.3


Organic bases: only one main kind- nitrogen atom

H3C NHH

H-Cl + H3C NHH

HCl-

N

Pyridine

(CH3CH2)3N

Triethylamine

H3CO

H H3C CH3

O

Oxygen-containing compounds: act as bases with strong acidsalso act as acids (O-H, C-H) with bases


2.11 Acids and Bases: The Lewis Definition

Lewis Acid: electron pair acceptorLewis Base: electron pair donor

B A+ B A

filledorbital

vacantorbital

Lewis base Lewis acid


Cl H O HH

+ O HH

H+ Cl-

Lewis Acid / Base Reaction

BF

FF

O CH3CH3

+ BF

FF

OCH3

CH3

AlCl

ClCl

NCH3

CH3+ AlCl

ClCl

NCH3

CH3CH3

CH3


Lewis acids:

neutral proton donors: H2O HCl HBr HNO3 H2SO4

CH3COOH CH3CH2OHOH

cations: Li+ Mg2+ Br+

metal compounds: AlCl3 BF3 TiCl4 FeCl3 ZnCl2 SnCl4

• Lewis definition of acid/base is broader than BrØnsted-Lowry definition: Lewis acid include many other species other than H+


AlCl3 BF3tetrahedral complex

(maximum 4 CN)

Common Lewis Acids

OBBr3O

AlCl3


Cl H O HH

+ O HH

H+ Cl-

Lewis base

Lewis Base: nonbonding electron pair


Lewis bases:

CH3CH2OH CH3OCH3 CO

H3C H CO

H3C CH3

CO

H3C Cl CO

H3C OH CO

H3C OCH3 CO

H3C NH2

CH3SCH3NCH3

CH3H3C

• most oxygen and nitrogen containing compounds are Lewis bases


CH3CH2OH CO

H3C OH

Some compounds, such as alcohols and carboxylic acids, can act as both acids and bases.

CH3CH2O CO

H3C O HHBrØnsted-Lowry Acid: proton (H+) donor

Lewis base: electron pair donor


H3C CH3

O+ H2SO4

H3C CH3

O+ BF3

H3C CH3

O+

H

HSO4-

H3C CH3

OBF3


H3C O

O+ H2SO4

H

- Reaction normally occurs only once in such instances, and the more stable of the two possible protonation products is formed.

• Some compounds, such as carboxylic acids, ester, and amides, have more than one atom with a lone pair of electrons and therefore react at more than one site.

For acetic acid, protonation occurs on the doubly bonded oxygen.

more stable(resonance stabilization) H3C O

O

H

H

+ HSO4-

H3C O

O

H

H

or H3C O

O

H

H


CH

CH

HCH

H

HCH

HH

CH HH

CH3CH2CHCH3 CH3CH2CH(CH3)2

CH3or

condensed structure

2.12 Drawing Chemical Structures

condensed structure: C-H and C-C single bonds aren't shown


H2CC

CH

CH2

CH3

H2CH2C

CH2

CH2

CH

H2C CH3 CH3

skeleton structure: line drawing of C-C bond- carbon atoms aren't usually shown- hydrogen atoms bonded to carbons aren't shown- atoms other than C, H are shown- C, H can be indicated for emphasis and clarity


CC

CCC

CH

H

HOH

H

HOH

NHCH3HO

HO

OH

Adrenaline

HNHO

HO

OH

CH3

Space-filling model

Ball-and-stick modelWire-frame model

Molecular Models

Buckminsterfullerene (C60)

Alkaloids: Naturally Occurring BasesChemistry @ Work

alkaloid: amines derived from natural sources

C NH

H

HH

HMethylamine (found in rotting fish)

O

HO

HO

N CH3

Morphine

NCH3

OCH2OH

O

Atropine

CH3

OH

HNCH3

Ephedrine

polar covalent bonds: electron distribution is...

Documents

Transcript of polar covalent bonds: electron distribution is...