What’s coming up???• Oct 25 The atmosphere, part 1 Ch. 8• Oct 27 Midterm … No lecture• Oct 29 The atmosphere, part 2 Ch. 8• Nov 1 Light, blackbodies, Bohr Ch. 9• Nov 3,5 Postulates of QM, p-in-a-box Ch. 9• Nov 8,10Hydrogen and multi – e atoms Ch. 9• Nov 12 Multi-electron atoms Ch.9,10• Nov 15 Periodic properties Ch. 10• Nov 17 Periodic properties Ch. 10• Nov 19 Valence-bond; Lewis structures Ch. 11• Nov 22 VSEPR Ch. 11• Nov 24 Hybrid orbitals; VSEPR Ch. 11, 12• Nov 26 MO theory Ch. 12• Nov 29 MO theory Ch. 12• Dec 1 bonding wrapup Ch. 11,12• Dec 2 Review for exam

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ATOMIC NUMBER

-ELE

CT

RO

N A

FF

INIT

Y

H

He

Li

Be

B

C

N

O

F

Ne

Na

Mg

Al

Si

P

S

Cl

Ar

Atoms combine to form compounds in an

attempt to obtain a stable noble gas electron

configuration with 8 electrons in the valence

shell

THE OCTET RULE

A stable electronic configuration can be attained in two ways…….

ELECTRON SHARING

ELECTRON TRANSFER IONIC BONDING

COVALENT BONDING

LEWIS MODEL OF BONDING

electron transferIONIC BONDING

Lewis Symbol symbol for elementand a dot for each

valence electron

1s22s22p6 3s1 1s22s22p63s23p5

ClNa

Na+1s22s22p6 1s22s22p63s23p6

ClNa Cl

electron sharing

F F

COVALENT BONDING

Atoms go as far as possible toward

completing their octets by sharing electron pairs

Lewis Symbols…..

Consider F2 F FThe electronic configuration of F is 1s22s22p5

These are combined to form FThese are combined to form F22

electron sharing

F 1s22s22p5 F 1s22s22p5+

F F+ F F or F F

non-bonding, or lone pair of electrons

bonding pair of electrons

COVALENT BONDING

Atoms go as far as possible toward

completing their octets by sharing electron pairs

The simplest molecule

The type of bond that forms depends on theelectronegativity difference between the two

atoms involved in the bond

IONIC OR COVALENT?

IONIC VERSUS COVALENT BONDS

Compounds composed of elements with a large difference in

ELECTRONEGATIVITY tend to have

significant ionic character in their bonding

A B

B has a greater share

H +

+

F FH

HYDROGEN FLUORIDE

Fluorine is more electronegative than hydrogen.

LEWIS BONDING MODEL

Lewis structures are based on a

localized electron model

Electrons are always localized in one of two ways:

As lone pairs on a specific atom

OR

As bonding pairs between two specific atoms

Drawing Lewis Structures

O HH

non-bonding, or lone pair of electrons

bonding pair of electrons

Lewis structures consider only valence electrons

Water

How do we draw these?????

H2O

Building Lewis (ELECTRON DOT) Structures of Molecules.

HCN as an example...

Step 1. Count the total number of valence electrons

C has 4 N has 5 H has 1 Total of 10

Step 2. Place one e- pair between each BONDED atomH C N

We have 6 e- left All atoms must have an octet or duet

Add electrons to terminal atoms firstStep 3.to get an octet or duet.

Building Lewis (ELECTRON DOT) Structures of Molecules.

Add 6 electrons in pairs to give the N an octet.

Step 3. Add remaining electrons to terminal atoms first

H C N

Step 4. Add any electrons left over to central atom

Do all atoms have an octet?

We have none left!

Step 5. Check for an acceptable Lewis Structure

IN THIS CASE

Building Lewis (ELECTRON DOT) Structures of Molecules.

H C N

Step 5. Check for an acceptable Lewis Structure

bring electron pairs from outer N atom to form shared pairs to give C its octet!!!

H C N

Still no octet on C Do it again!!!!

H C N

three electron pairs between the C and N………...

Another possible structure is….

H C N

Building Lewis (ELECTRON DOT) Structures of Molecules.

H C N

H N C

How can we choose?

FORMAL CHARGE

Structures with the lowest formal charges are likely to

have the lowest energy.

Each atom is assigned all of its lone electrons and

Formal charge =

#valence electrons

_ #unshared electrons

_ 1/2#shared electrons{ } { }

gives an indication of the extent to which atoms

FORMAL CHARGE

have gained or lost electrons in the process of covalent

bond formation.

half of the electrons bonded to it.

All possible Lewis structures with stable

electronic configurations for HCN and HNC.

Calculate formal charge for this one

FC on C = 4 - 0 - 1/2 (8) = 0

_ 1/2#shared electrons{ }Formal

charge =#valence electrons

_ #unshared electrons

{ }

FC on N = 5 - 2 - 1/2 (6) = 0

Hydrogen is zero

H C N

..

H N CH C N

All possible Lewis structures with stable

electronic configurations for HCN and HNC.

0 0 0 0 +1 -1

WE CHOOSE THE STRUCTURE WITH THE

FORMAL CHARGES CLOSEST TO ZERO

AND ANY NEGATIVE FORMAL CHARGES ON THE MOST ELECTRONEGATIVE ELEMENTS

H N CH C N

Example: H2CO

Total number of valence electrons = 6+4+1+1=12

Put carbon in the middle…..

Place electrons between atoms

Now add remaining electrons to O

Now O has an octet but C does not….

Share a pair between C and O

Now we have a double bond between C and O

H

H

OC

Example: H2CO

Total number of valence electrons = 6+4+1+1=12

Give C an octet

We can write this….

H

H

OC

H

H

OC

C

H

H

O

N has five 1s22s22p3

O has six 1s22s22p4

Count up valence electrons

Let’s look at the nitrate anion NO3-

Plus one extra for negative charge

Valence electrons = 5 + 3 x 6 + 1 = 24

Put a pair between each atomNitrate anion NO3

-

N

O

O O

nitrogen does not have noble gas structure!!!

form a double bond by sharing a pair from one of the oxygen atoms……….

FORM A DOUBLE BOND BETWEEN O AND N

Here is one

Here is another!Here is another!

N

O

O O

-

N

O

O O

-

N

O

O O

-

Experiment shows all three bonds are the same.

Any one of the structures suggests one is different!

O O

O

N

All bond angles 120 0

All bond lengths 128 pm

N

O

O O

- Double Bond

Single Bond

Should be different!

So…….

We use a double headed arrow between the structures..

The electrons involved are said to be DELOCALIZED over the structure.

The blended structure is a RESONANCE HYBRID

RESONANCE

N

O O

O

O O

O

N N

O O

O

S

O O

S

O O

SO2

We say that the real SO2 molecule is a hybrid of the

two resonance forms.

Experiment shows that both S-O bonds are

equivalent.

LEWIS STRUCTURE

EXCEPTIONS TO THE OCTET RULE…….

Molecules with more than 8 electrons around central atom.

Molecules with unpaired electrons.

Molecules with less than an octet around central atom

Lets do SF6……..

Elements in rows 3 and following can exceed the octet rule:

F

S

F

F

F

F

F

SF6

I II

-I3

When it is necessary to exceed the octet rule the extra electrons go on the central third row element.

S … 12

Central I … 10

Molecules which have unpaired electrons.

NO2Is a free radical

Total number of valence electrons = 5+6+6 = 17

O N O

FREE RADICALS

Form double bond to get N close to octet

O N O O N O

RESONANCE