Chapter 1 5 Chem207

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Chem 207 B. R. Kaafarani 1

Chapter 1Structure Determines Properties

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Atomic number (Z) = number of protons in nucleus.

(this must also equal the number of electrons

in neutral atom)

Mass number (A) = sum of number of protons

+ neutrons in nucleus.

XZ

A

1.1. Atomic Number and Mass Number

C6

12

O8

16

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Schrdinger Equation Schrdinger combined the idea that an electron has

wave properties with classical equations of wavemotion to give a wave equation for the energy of an

electron in an atom.

Wave equation (Schrdinger equation) gives a

series of solutions called wave functions ().

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Wave Functions Only certain values of are allowed.

Each corresponds to a certain energy.

The probability of finding an electron at a particular

point with respect to the nucleus is given by 2.

Heisenberg uncertainty principle: we can not tell

where an electron is; however, we can predict where it is

most likely to be. Each energy state corresponds to an orbital.

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A boundary surface encloses the regionwhere the probability of finding an electronis highon the order of 90-95%

Boundary Surface

Figure 1.2. Boundary surfaces of a 1s orbital and a 2s orbital.

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Quantum Numbers Each orbital is characterized by a unique set of

quantum numbers.

The principal quantum numbern is a whole number

(integer) that specifies the shell and is related to the

energy of the orbital.

The angular momentum quantum numberl is usually

designated by a letter (s,p,d,f, etc) and describes the

shape of the orbital.

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Quantum Numbers Each orbital is characterized by a unique

set of quantum numbers. The magnetic quantum numberml describes the

orientation of the orbital in space.

The electron spin quantum numberms has a value

of +1/2 or -1/2.

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s Orbitals

s Orbitals are spherically symmetric.

The energy of an s orbital increases with thenumber ofnodal surfaces it has.

A nodal surface is a region where the probability of

finding an electron iszero.

A 1s orbital has no nodes; a 2s orbital has one;

a 3s orbital has two, etc.

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The Pauli Exclusion Principle

No two electrons in the same atom can have the

same set of four quantum numbers.

Two electrons can occupy the same orbital only

when they have opposite spins.

There is a maximum of two electrons per orbital.

1s 2s 2p

C 6

Z

Hunds rule: es fill orbital of equal energy first.

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Spherical Dumbbell- shaped along X, Y, Z axis

8O: 1s22s22p4

s: 1 orbital, 2 electronsp: 3 orbitals, 6 electrons

d: 5 orbitals, 10 electronsf: 7 orbitals, 14 electrons

11Na: 1s22s22p63s1

x

y

z

s px

py

pz

node

3Li: 1s22s1

n indicates the row the element is in.

Valence e-s

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Octet rule: Nobel or rare gases (helium, neon, andargon) have 8 electrons in their valence shell.

Ionization energy is the energy required to removean e- from an atom.

Electron affinity is the energy change upon the

addition of an e- to an atom.

Energy absorbing reactions are calledendothermic.

Energy releasing reactions are calledexothermic.

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Cl (anion)

Na+ (cation)

1.2. Ionic Bonding

An ionic bond is the force of electrostatic attraction

between oppositely charged ions.

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Ionic Bonds Ionic bond is a bond between atoms of opposite charges;these atoms are called ions. +ve ions are called cations; -ve ions are called anions.

Atoms lose/gain electrons to become ions that have thesame configuration as the nearest noble gas.

1s22s22p63s1

Na (g) Na+ + e-

Sodium atom1s

2

2s2

2p6

Sodium ion

1s22s22p63s23p6

Cl-

Chloride ion

1s22s22p63s23p5

Cl (g) + e-

Chlorine atom

Na (g) Cl (g) Na+Cl-+

Ionic: electrostatic/coulombic

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1.3. Covalent Bonds

Gilbert Newton Lewis suggested Covalent Bonds- sharede- pair.

H2H

.Hydrogen atom .H

Hydrogen atom

H:HHydrogen molecule

Covalent bond

Lewis structures: e-s represented as dots

Bond dissociation energy: energy required to dissociate 2 Hs (435 kJ/mol)

F2 F F+ F F

Each fluorine has 8 e-

s in its valence shell. What rule?

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Methane has 1 C and 4 Hs

C H.... .+ 4 H:C:H

:

:

H

H

C HH

H

H

CH4

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Ethylene C::C

:

: :

:H

H

H

H C C

H

H H

H H H

H H

Acetylene H:C:::C:H H C C H H H

1.4. Double Bonds and Triple Bonds

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1.5. Polar Covalent Bonds and Electronegativity

In a bond between two atoms, when one atom has larger

tendency to attract e-s, the e- distribution is polarized. Polar

Covalent Bond. +HF-

Direct. of polar.

Electronegativity is the tendency of an atom to attract e

-

s toitself.

- Electronegative E- and electropositive E+ atoms.

E- increases across a row from left to right and decreases ingoing down a column.

Fluorine has the highest E- value 4.0 (Linus Pauling Scale).

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Pauling Electronegativity Scale

1.0

Na

0.9

Li Be B C N O F

1.5

Mg

1.2

2.0

Al

1.5

2.5

Si

1.8

3.0

P

2.1

3.5

S

2.5

4.0

Cl

3.0

Electronegativity increases from left to right in theperiodic table.

Electronegativity decreases going down a group.

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Generalization

Nonpolar bonds connect atoms of the same electronegativity.

HH : N N:F:..

..F:..

..

Polar bonds connect atoms of different electronegativity.

F:..

..H

O..

..H

H

: O C

O:.. ..

The greater the difference in electronegativity

between two bonded atoms; the more polar the bond.

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Red is negative charge;blue is positive.

F:..

..H

Solidsurface

Electrostatic potential maps show the chargedistribution within a molecule.

Transparentsurface

Electrostatic Potential Maps (EPM)

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Dipole Moment ()

= e x d

Charge on an electron 4.80 x 10-10 esu

Distances fall into 10-8 cm

debye, D = 10-18 esucm

H-F = 1.710-18 esu.cm = 1.7 D

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1.6. Formal Charges

NO

OO

H -+

:

:

:::

:

:

e- count (H)= (2) =1

e- count (O)= (4) + 4 =6 e- count (O)= (2) +6 =7

e- count (N)= (8) =4

e- count (O)= (4) + 4 =6

Formal charge = valence e-s e- count

or

Formal charge = group # in # of bonds # of unshared e-s

periodic table

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Constitutional Isomers Isomers are different compounds that have the samemolecular formula.

Constitutional or structural isomers differ in the way their

atoms are bonded.

Constitutionalisomers of C3H6

Constitutional

isomers of CH3NO2

Cyclopropane1-Propene

NOC

O

H

: :

:

+

H

H:: _

Nitromethane

:OC

H

H

H N O

: ::

:

Methyl nitrite

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1.8. Resonance

XX

OO O:

::

_

1.211.47:

:

:

+

OO O

:

:

:

_

:

:

:

+O

O O:

:

: _

:

:

+

Ozone:O3

Both bonds were equal in length: 1.28

OO O

:

:

+

:

:

:

-1/2-1/2

More than Lewis structure Hybrid structure

Lewis structures: localized.

Resonance Structures: delocalized.

1 pm = 10-12 m

1 = 10-10 m

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Table 1.6. Rules of Resonance

The positions of atoms stay the same, only the e-s positions are different.

NO

OH3C

+_ O NH3C O

nitromethane methyl nitrite

NO

O

H3C

Second row elements can not have more than 8 valence e-s (unstable).

O NH3C O O NH3C O+

The more stable resonance structure is the one that has the smallest # ofopposite charges.

When one atom bears a formal charge, the most stable resonancestructure is the one where the charge resides on the more E- atom.

N C O N C O:

:

::

::

: :

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Table 1.6. Rules of Resonance (contd)

The resonance structures must have the same net charge.

e- delocalization stabilizes a molecule. The degree of stabilization

is greatest when resonance structures are of equal stability.

O NH3C O O NH3C O+

_

N

O

H3C

O

+_ N

O

H3C

O

+

_

The resonance structures must have the same # of unpaired e-s.

NOH3C

O: :

:

+ :: _ NO

H3C:

:.

. : ::

O

NO

H3CO

: :

:

+:_ N

OH3C

:

:_: :

::O

:_

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Valence Shell Electron PairRepulsions (VSEPR)

1.10. The Shapes of Some Simple

Molecules

The most stable arrangement of groupsattached to a central atom is the one that has the

maximum separation of electron pairs (bonded or

nonbonded).

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Tetrahedral geometryHCH angle = 109.5

Table 1.7. Methane

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Bent geometryHOH angle = 105

But notice the tetrahedral arrangementof electron pairs.

O

H

..

H

:

Table 1.7. Water

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Trigonal pyramidal geometryHNH angle = 107

But notice the tetrahedral arrangementof electron pairs.

N

H

H

H

:

Table 1.7. Ammonia

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FBF angle = 120

Trigonal planar geometryallows for maximum separation

of three electron pairs.

Table 1.7. Boron Trifluoride

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Multiple Bonds

Four-electron double bonds and six-electron

triple bonds are considered to be similar to a two-electron single bond in terms of their spatial

requirements.

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HCH and HCOangles are close to 120

trigonal planar geometry

C OH

H

Table 1.7. Formaldehyde

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OCO angle = 180linear geometry

O C O

Table 1.7. Carbon Dioxide

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1.11. Molecular Dipole Moments

Molecular dipole moment is the result of all theindividual bond dipole moments of a substance.

C OO

C Cl

Cl

Cl

Cl

C Cl

Cl

HH

= 0 D

= 0 D

= 1.62 D

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Arrhenius

An acid ionizes in water to give protons. A base

ionizes in water to give hydroxide ions.

Brnsted-LowryAn acid is a proton donor. A base is a proton

acceptor.

LewisAn acid is an electron pair acceptor. A base is

an electron pair donor.

1.13. Acids & Bases

Definitions

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Arrhenius Acids and Bases

An acid is a substance that ionizes to give protons

when dissolved in water.

AH+

H A + ..

A base is a substance that ionizes to give hydroxideions when dissolved in water.

M+ + OH.. ....M OH..

..

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Arrhenius Acids and Bases

Strong acids dissociate completely in water.

Weak acids dissociate only partially.

AH+

H A + ..

Strong bases dissociate completely in water.

Weak bases dissociate only partially.

M+ + OH..

..

..

M OH....

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Acid Strength is Measured by pKa

Ka =[H+][A]

[HA]

pKa = log10Ka

AH+

H A + ..

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Brnsted-Lowry definition:

An acid is a proton donor

A base is a proton acceptor

1.14. Acids and Bases:The Brnsted-Lowry View

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H AB.. B H A..

+

A Brnsted Acid-Base Reaction

A proton is transferred from the acid to the

base.

+ +

base acid conjugateacid

conjugatebase

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H3O+: hydronium ion

H BrO

H

H

.. ..

H

Br

.. ....

....

....

H

..O H+

Proton Transfer from HBr to Water

base acid conjugate conjugate

acid base

++

WithBrnsted-Lowry View : Acid does not dissociatein H2O: it transfers a proton to water. Therefore, wateracts as a base.

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H OHN

H

H

.. ..

H

H

..N H OH........

..

Water as a Brnsted Acid

base acid conjugate conjugateacid base

+ +

Dissociation Constants (pK ) of Acids

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Weakeracid

Strongeracid

HI I

HBr Br

HCl Cl

H3O+

-10.4

-5.8

-4.8

-3.9-1.7 H2O

H2SO4 HSO4

Dissociation Constants (pKa) of Acids

Acid pKa Conj. base

Strong acids are stronger than hydronium ion.

The stronger the acid, the weaker the conjugate base.

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Weak acids are weaker than hydronium ion.

H3O+ 1.7 H2O

HF 3.5 F

CH3CO2H 4.6 CH3CO2

NH4+ 9.2 NH3

H2O 15.7 HO


Acid pKa Conj. base

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CH3OH CH3O

CH3CH2OH ~16 CH3CH2O

(CH3)2CHOH ~17 (CH3)2CHO

(CH3)3COH ~18 (CH3)3CO

H2O 15.7 HO

15.2


Acid pKa Conj. base

Alcohols resemble water in acidity; their conjugatebases are comparable to hydroxide ion in basicity.

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NH3 ~36 NH2

(CH3)2NH ~36 (CH3)2N


Acid pKa Conj. base

Ammonia and amines are very weak acids;their conjugate bases are very strong bases.

Dissociation Constants (pK ) of Acids

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26HC CH

43

45

62CH3CH3

H2C CH2

H

HH

H

H H

HC C

HH

H

H H

H2C CH

CH3CH2


Acid pKa Conj. base

Most hydrocarbons are extremely weak acids.

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1.15. What Happened to pKb?

A separate basicity constantKb is not necessary.

Because of the conjugate relationships in theBrnsted-Lowry approach, we can examine acid-

base reactions by relying exclusively on pKa values.

Example H

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Example

N

H

H

HN

Which is the stronger base, ammonia (left) or

pyridine (right)?

Recall that the stronger the acid, the weaker the

conjugate base. Therefore, the stronger base is the

conjugate of the weaker acid. Look up the pKa values of the conjugate acids of

ammonia and pyridine in Table 1.8.

Example

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Example

pKa = 9.3

pKa = 5.2

weaker acid

stronger acid

Therefore, ammonia is astronger base than pyridine.

N

H

N

H

HHH

1 16 How Structure Affects Acid Strength

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The strength of the bond to the atom from which

the proton is lost.

The electronegativity of the atom from which the

proton is lost.

Changes in electron delocalization on ionization.

1.16. How Structure Affects Acid Strength

Bond Strength

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Bond Strength

Bond strength is controlling factor when comparing

acidity of hydrogen halides.

HF HCl HBr HIpKa 3.1 -3.9 -5.8 -10.4

weakest acid strongest acid

strongest HX bond weakest HX bond

Bond Strength

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Bond Strength

Recall that bond strength decreases in a group in

going down the periodic table.

Generalization: Bond strength is most importantfactor when considering acidity of protons bonded to

atoms in same group of periodic table (as in HF, HCl,

HBr, and HI).

Another example: H2S (pKa = 7.0) is a stronger acid

than H2O (pKa = 15.7).

Electronegativity

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Electronegativity

Electronegativity is controlling factor when

comparing acidity of protons bonded to atoms inthe same row of the periodic table.

Electronegativity

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Electronegativity

pKa

60 36 15.7 3.1

weakest acidstrongest acid

least electronegative most electronegative

CH4 NH3 H2O HF

Electronegativity

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Electronegativity

H AO

R

H

.. ..

R

H

..O H A..

+

++

The equilibrium becomes more favorable as Abecomes better able to bear a negative charge.

Another way of looking at it is that H becomes morepositive as the atom to which it is attached becomes

more electronegative.

Bond Strength Versus Electronegativity

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Bond strength is more important when comparing

acids in which the proton that is lost is bonded to

atoms in the same group of the periodic table.

Electronegativity is more important when comparing

acids in which the proton that is lost is bonded to

atoms in the same row of the periodic table.

Bond Strength Versus Electronegativity

Acidity of Alcohols

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Acidity of Alcohols

pKa

15.7

15.2

16

17

18

Alcohols (ROH)

resemble water

(HOH) in theiracidity.

HOH

CH3OH

CH3CH2OH

(CH3)2CHOH

(CH3)3COH

In many acids theacidic proton is

bonded tooxygen.

Acidity of Alcohols

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y

CH3CH2OH CF3CH2OH16 11.3

weaker stronger pKa

Electronegative substituents can increase the acidity

of alcohols by drawing electrons away from the OH

group.

Inductive Effect

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C C O H

H

H

F

F

F

+

The greater acidity of CF3CH2OH compared to

CH3CH2OH is an example of an inductive effect.

Inductive effects arise by polarization of theelectron distribution in the bonds between atoms.

Electrostatic Potential Maps

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p

CH3CH2OH CF3CH2OH

The greater positive character of the proton of theOH group of CF3CH2OH compared to CH3CH2OH isapparent in the more blue color in its electrostatic

potential map.

Another example of the inductive effect

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O

CH3C O H

O

CF3C O H

4.7 0.50

weaker stronger

pKa

p

Electron Delocalization

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Ionization becomes more favorable if electrondelocalization increases in going from left to right in

the equation.

Resonance is a convenient way to show electron

delocalization.

H AO

R

H

.. .. A..

R

H

..O H+

++

Nitric Acid

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O

+

O NH

O

O

H

H

HO

H

H

+

O

+

O N

O

+

pKa = -1.4

+

Nitrate ion is stabilized

by electron delocalization.

Nitric Acid

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O

+

O N

O

O

+

O N

O

O

+

O N

O

Nitrate ion is stabilized by

electron delocalization.

Negative charge is sharedequally by all three oxygens.

Acetic Acid

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O

O CH

O

H

H

pKa = 4.7

+

CH3

HO

H

H

+

+

O

O C

CH3

Acetic Acid

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OO C

OO C

CH3 CH3

Acetate ion is stabilized by electron delocalization.

Negative charge is shared equally by both oxygens.

1.17. Acid-Base Equilibria

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Generalization

Stronger acid + Stronger base Weaker acid + Weaker base

The equilibrium in an acid-base reaction is favorableif the stronger acid is on the left and the weaker acid is

on the right.

Example of a strong acid

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H

H

..O H Br

.. ....

..

+H BrO

H

H

.. ....

....

pKa = -5.8stronger acid pKa = -1.7weaker acid

+ +

The equilibrium lies to the side of the weaker acid(to the right).

Example of a weak acid

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pKa = 4.7weaker acid

pKa = -1.7stronger acid

HOCCH3

O

+O

H

H

OH

H

H

+ + OCCH3

O

The equilibrium lies to the side of the weaker acid(to the left).

1.18. Lewis Acids and Lewis Bases

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Arrhenius

- An acid ionizes in water to give protons. A baseionizes in water to give hydroxide ions.

Brnsted-Lowry

- An acid is a proton donor. A base is a protonacceptor.

Lewis

- An acid is an electron pair acceptor. A base is anelectron pair donor.

Definitions

Lewis Acid-Lewis Base Reactions

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The Lewis acid and the Lewis base can be either

a neutral molecule or an ion.

Lewis acid Lewis base+

A + AB+

A+

ABB

+ +

A + ABB

A + ABB

+

B

Example: Two Neutral Molecules

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F3B + O

CH2CH3

CH2CH3

+F3B O

CH2CH3

CH2CH3

Lewis acid Lewis base

Product is a stable substance. It is a liquid with a boiling

point of 126 C. Of the two reactants, BF3 is a gas and

CH3CH2OCH2CH3 is a liquid with a boiling point of 34 C.

Boron trifluoride

Chapter 1 5 Chem207

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