chapter 8 basic concepts in chemical bonding. introduction attractive forces that hold atoms...

CHAPTER 8

Basic Concepts in Chemical Bonding

Introduction

Attractive forces that hold atoms together in compounds are called chemical bonds. The electrons involved in bonding are usually those in the outermost (valence) shell.

Introduction

Chemical bonds are classified into two types:

Ionic bonding results from electrostatic attractions among ions, which are formed by the transfer of one or more electrons from one atom to another.

Covalent bonding results from sharing one or more electron pairs between two atoms.

Lewis Dot Representations

of AtomsThe Lewis dot representation (or Lewis dot formulas) convenient bookkeeping method for

valence electrons electrons that are transferred or

involved in chemical bonding


of Atoms

Li Be B C N O F Ne

H.

He


of Atoms

elements in same group have same Lewis dot structures

Li Be B C N O F Ne.... .. ..

..HeH

.

.. . .

.. ....

...

..

.. .. .. .... ..

.


of Atoms

elements in same group have same Lewis dot structures

Li Be B C N O F Ne.... .. ..

..HeH

.

.. . .

.. ....

...

..

.. .. .. .... ..

.

Li & Na. .

N & P.. ..

..

. ..

. F & Cl...

....

.

... ..

.

Ionic Bonding

We can also use Lewis formulas to represent the neutral atoms and the ions they form.

Li + F.....

.. . Li+ F[ ]..

.... ..

Ionic Bonding

Li+ ions contain two electrons same number as heliumF- ions contain ten electrons same number as neonLi+ ions are isoelectronic with heliumF- ions are isoelectronic with neon Isoelectronic species contain the same number of electrons.

Ionic Bonding

potassium reacts with bromine equation for the reaction

Ionic Bonding

potassium reacts with bromine2 K(s) + Br2 (l) + Br-(s)

electronically this is occurring 4s 4pK [Ar] Br [Ar] becomes

Ionic Bonding

4s 4pK+ [Ar]Br- [Kr]Lewis dot formulas

Ionic Bonding

cations become isoelectronic with preceding noble gasanions become isoelectronic with following noble gas

K + Br.....

.. . K+ Br[ ]..

.... ..

Covalent Bondingcovalent bonds form when atoms share electronsshare 2 electrons - single covalent bondshare 4 electrons - double covalent bondshare 6 electrons - triple covalent bondattraction is electrostatic in nature lower potential energy when bound

Covalent Bonding

Covalent Bonding

Lewis dot representation H molecule formation

HCl molecule formation+H. H . H H.. or H2

H Cl H Cl+... ..

.. ....

..

... or HCl

Covalent Bonding

extremes in bondingpure covalent bonds electrons equally shared by the atomspure ionic bonds electrons are completely lost or

gained by one of the atoms most compounds fall somewhere between these two extremes

Lewis Dot Formulas for Molecules and Polyatomic Ions

homonuclear diatomic molecules hydrogen, H2

fluorine, F2

nitrogen, N2

H HorH H..

F F.. .. ....

..

.. .. F F.. .... ..

.. ..or

N N········ ·· N N·· ··or


heteronuclear diatomic moleculeshydrogen halides hydrogen fluoride, HF

hydrogen chloride, HCl

hydrogen bromide, HBr

or ··H F··

··H F..

·· ····

or ··H Cl··

··H Cl..

·· ····

or ··H Br··

··H Br..

·· ····


water, H2O

H

H

O···· ··

··


ammonia molecule , NH3

H

H

N···· ··

·· H


ammonium ion , NH4+

H

H

N···· ··

·· H

H +

The Octet Rule

Elements try to a achieve noble gas configuration in most of their compounds.Lewis dot formulas are based on the octet rule.

N - A = S rule

N = # of electrons needed to be noble gas usually 8 2 for HA = # of electrons available in outer shells equal to group # 8 for noble gases

N - A = S rule

for ions add one e- for each negative charge subtract one e- for each positive chargecentral atom in a molecule or polyatomic ion is determined by: atom that requires largest number of

electrons for two atoms in same group - less

electronegative element is central

Drawing Lewis Dot Formulas

Write Lewis dot and dash formulas for hydrogen cyanide, HCN.

N = 2 + 8 + 8 = 18A = 1 + 4 + 5 = 10S = 8

H C N·· ·· ···· H C N ··or··


Write Lewis dot and dash formulas for the sulfite ion, SO3

2-. N = 8 + 3(8) = 32A = 6 + 3(6) + 2 = 26S = 6

O S O

O··

····

····

··

··

··

····

··

····

2-O S

O

O·· ·· ·· ···· ·· ··

······

2-or


Write Lewis dot and dash formulas for sulfur trioxide, SO3.

N = 8 + 3(8) = 32A = 6 + 3(6) = 24S =8

orO S O

O··

····

····

··

····

····

·· ·· O S

O

O·· ···· ·· ··

······

Resonance

three possible structures for SO3

two or more dot structures necessary to show bondinginvoke resonance Double-headed arrows are used to indicate

resonance formulas.

O S

O

O·· ···· ·· ··

······

OS

O

O·· ···· ·· ··

··

······

O S

O

O·· ···· ·· ··

····

Resonance Structures

Some molecules are not described by Lewis Structures very well.

Typically, structures with multiple bonds can have similar structures with the multiple bonds between different pairs of atoms.

Resonance

flaw in our representations of moleculesno single or double bonds in SO3

all bonds are the samebest picture

incorrect shape

OSO

O

--- ---

Limitations of the Octet Rule for Lewis Formulas

species that violate the octet rule N - A = S rule does not apply

covalent compounds of Be, B, Ncovalent compounds with expanded octetsspecies containing an odd number of electrons


In cases where the octet rule does not apply, the substituents attached to the central atom nearly always attain noble gas configurations. The central atom does not.


Write dot and dash formulas for BBr3.

B··. Br··

··

··.

BBr Br

Br

········

····

····

····

····

Br B

Br

Br····

··

·· ····

·· ····

or


Write dot and dash formulas for AsF5.

As··..

. F····

··.

··

As

F

F F

F F

····

··

·· ······

······

····

·· ··or

····

····

··

·· ······

······

···· AsF

F F

FF

······ ··

··

··

Formal Charges

It is possible to draw more than one Lewis structure with the octet rule obeyed for all the atoms.To determine which structure is most reasonable, we use formal charge.

Formal Charges

The charge on an atom that it would have if all the atoms had the same electronegativity.To calculate formal charge, electrons are assigned as follows:

•All nonbonding electrons are assigned to the atom on which they are found.•Half the bonding electrons are assigned to each atom in a bond.

Formal Charges

valence electrons – (½ number of bonds + lone pair electrons)

Consider:

C N

Formal Charges

For C: There are 4 valence electrons.In the Lewis structure there are 2 nonbonding electrons and 3 from the triple bond. There are 5 electrons from the Lewis structure.

Formal charge: 4 - 5 = -1

C N

Formal ChargesConsider:Consider:

For For NN::There are 5 valence electrons.There are 5 valence electrons.In the Lewis structure there are 2 nonbonding In the Lewis structure there are 2 nonbonding electrons and 3 from the triple bond. There electrons and 3 from the triple bond. There are 5 electrons from the Lewis structure.are 5 electrons from the Lewis structure.

Formal charge = 5 - 5 = 0Formal charge = 5 - 5 = 0We write:We write:

C N

C N

Formal Charges

The most stable structure has:•the smallest formal charge on each atom,•the most negative formal charge on the most electronegative atoms.

Polar and Nonpolar Covalent Bonds

polar covalent bonds unequally shared electrons assymmetrical charge distribution different electronegativities

ElectronegativityElectronegativity


bondpolar very 1.9 Difference

4.0 2.1 ativitiesElectroneg F H

1.9


Electron density map of HF blue areas - low

electron density red areas - high

electron densitypolar molecules have separation of centers of negative and positive charge


bondpolar slightly 0.4 Difference

2.5 2.1 ativitiesElectronegI H

0.4


Electron density map of HI blue areas - low

electron density red areas - high

electron densitynotice that the charge separation is not as big as for HF HI is only slightly polar

Continuous Range of Bonding Types

all bonds have some ionic and some covalent character HI is about 17% ionicgreater the electronegativity differences the more polar the bond

Continuous Range of Bonding Types

molecules whose centers of positive and negative charge do not coincide are polar have a dipole momentdipole moment has the symbol is the product of the distance,d, separating charges of equal magnitude and opposite sign, and the magnitude of the charge, q

Dipole Moments

molecules that have a small separation of charge have small molecules that have a large separation of charge have a large example HF and HI

Dipole Moments

in molecules some nonpolar molecules have polar bonds

2 conditions must hold for a molecule to be polar

units Debye0.38 units Debye1.91 I- H F- H

--

Dipole Moments

There must be at least one polar bond present or one lone pair of electrons.

The polar bonds, if there are more than one, and lone pairs must be arranged so that their dipole moments do not cancel one another.


nonpolar covalent bonds electrons are shared equally symmetrical charge distributionmust be the same element to share exactly equallyH2

N2

H HorH H..

N N········ ·· N N·· ··or

Bond Energies

Bond energy - amount of energy required to break the bond and separate the atoms in the gas phase

A - B bond energy A + B

H - Cl H + Clg g g

gkJ

mol g g

432

Bond EnergiesTable of average bond energiesMoleculeBond Energy (kJ/mol)

F2 159 Cl2 243

Br2 192O2 (double bond) 498N2 (triple bond) 946

Bond EnergiesBond energies can be calculated from otherHo

298values

Bond Energies

Example: Calculate the bond energy for hydrogen fluoride, HF.

Bond Energies

Example: Calculate the bond energy for hydrogen fluoride, HF.

H - F BE H F atoms NOT ions

H - F H F H BE

H H H H

H kJ + 78.99 kJ kJ

H kJ BE for HF

g HF g g

g g g 298o

HF

298o

f Ho

f Fo

f HFo

298o

298o

g g g

or

218 0 271

568 0

.

.

Bond Energies

Example: Calculate the average N-H bond energy in ammonia, NH3.

Bond Energies

bonds H-N molkJ

H-N

o298

o298

oNH f

oH f

oN f

o298

H-No298ggg3

3913

kJ 1173BE average

kJ 1173H

kJ 11.46)218(3)7.472(H

HH 3HH

BE 3H H 3 + NNH

g3gg

Bond Energies

In gas phase reactions Ho values may be related to bond energies of all species in the reaction.

formed bondsbroken bonds DDH rxn

Bond Energies

Example: Use the bond energies listed in Table 8.4 to estimate the heat of reaction at 25oC for the reaction below.

CH O CO 2 H O

H BE BE BE 4 BE4 g 2 g 2 g 2 g

298o

C-H O=O C=O O-H

2

4 2 2

Bond EnergiesExample: Use the bond energies listed in Table 8.4 to estimate the heat of reaction at 25oC for the reaction below.

CH O CO 2 H O

H BE BE BE 4 BE

H kJ

H kJ

4 g 2 g 2 g 2 g

298o

C-H O=O C=O O-H

298o

298o

2

4 2 2

4 414 2 498 2 741 4 464

686

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chapter 8 basic concepts in chemical bonding. introduction attractive forces that hold atoms...

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