Chemistry 445Inorganic Chemistry

Lecture 1.

Lewis Dot StructuresVSEPR

Gilbert Newton Lewis (1875-1946)

G. N. Lewiswas probablythe best chemistwho never won the Nobel Prize

Lewis Dot Structures (revision)

Lewis dot structures present a simple approach to bonding that allows us to rationalize much molecular structure. The idea is that atoms share electrons in the valence shell to form the chemical bond, with one pair of electrons per bond. Note that each H-atom has two electrons, which is the structure of He, the next inert gas.

H-atom H-atom H2 molecule

Electron pair = single bondValence electrons

(Each H-atom has one valence electron)

Lewis Dot Structures (contd.):

O-atom O-atom O2 molecule

Periodic table

Two shared pairs of electrons= double bond

1 2 3 4 5 6 7 8

Oxygen has six valence electrons

The octet ruleElectrons are shared in forming bonds such that atoms have the same number of electrons in their valence shells as the nearest noble gas, including the electrons shared with the atom to which they are bonded.

O-atom O-atom O2 molecule

Each oxygen atom in the O2 molecule now has eightvalence electrons, including those it shares with theother oxygen atom = number of electrons (8 = octet)in the nearest inert gas = neon.

8.1 Chemical Bonds, Lewis Symbols, and

the Octet rule.Chemical bonding involves mainly the attempt to achieve the rare gas number of valence electrons, i.e. an octet. This can be achieved in several ways.

Ionic bond: Electrons are mainly the property of one of the two atoms forming the bond.

Covalent bond: Electrons are shared so that each atom has a noble gas electronic configuration.

Metallic bonds. Electrons are lost into the conduction band.

8.2 Ionic Bonding.

This occurs between metallic elements from the left-hand side of the periodic table and non-metallic elements from the right hand side of the periodic table.

Note that Na gives up its lone valence electron to Cl, so that they both end up with an octet of electrons.

8.3 Covalent bonding.

Here the two atoms share the electrons to achieve a covalent bond.

two pairs of electrons equally sharedbetween the two oxygen atoms

Multiple Bonds and bond order:The sharing of a single pair of electrons consititutes a single bond. Sharing of two pairs of electrons constitutes a double bond, and sharing three pairs of electrons constitutes a triple bond.

H:H :O::O: :N:::N:Single bond double bond triple bond

Bond order: a single bond has bond order = 1, a double bond has bond order = 2, and a triple bond has a bond order = 3. Fractional bond orders such as 1½ or 1⅓ are also possible, as discussed below.

.. ..

Some more examples of Lewis dot structures:

The N2 molecule:

N-atom N-atom N2 molecule

Periodic table

1 2 3 4 5 6 7 8

triple bond

Examples of Lewis dot diagrams:

Methane, CH4:

Carbon has four valence electrons (red)

H C

H

H

H

Carbon achievesoctet of electrons

Hydrogensachieve twoelectrons like He

One sharedpair of electrons= single bond

single line= single bond

Carbon dioxide: (CO2)

Examples of Lewis dot diagrams:

Carbon has four valence electrons (red)

oxygens have six valence electrons (black)

O=C=O

Carbon and both oxygensachieve an octet of electrons

two sharedpairs of electrons= double bond

double line =double bond

Examples of Lewis dot diagrams:

Sulfur dioxide: (SO2)

O=S-O(or O-S=O ?)

O S O

actual structure is averageof the two (bond order = 1½) :

doublebond?

singlebond?

SO2 is an example where a molecule can be written in two ways and actual structureis the average of the two. Thisis called RESONANCE (see later)

Slightly different Lewis dot representations:

One can also represent molecules/ions with a combination of dots and lines for bonds, remembering that each line represents a shared pair of electrons, e.g. the phosphate anion:

8.6 Resonance structures: Ozone (O3)

OO

O OO

O

.. ..

= OO O

The ozone molecule can be written with two equivalent Lewis dot structures. In such a situation the actual structure is the average of these two structures, with the two O-O bond lengths equal.

bond order = 1½

The ozone molecule

O-O bonds = 2.78 Å

double arrow= resonance

OO O

Resonance structures – the nitrite anion: (NO2-)

ON

O ON

ON

O O:

: :

::

..:

::

:

:..

=- - -

In drawing up a Lewis dot diagram, if we are dealing withan anion, we must put in an extra electron for each negative charge on the anion:

O

NO

:: :

::

.

:: :. -

negative chargeon anion

One extra electronin Lewis dotdiagram becauseof single negative charge on anion

Two resonance structures average structure

Bond order = 1½

The nitrate anion:

ON

O:

: :

::

..O

....

O

NO

::

:

::

..O.... O

N

O

::

:

: :

..O..

..

ON

O

O

- - -

-average bondorder (B.O.)=

2 + 1 + 1 = 1⅓ 3

B.O. = 2 B.O. = 1 B.O. = 1

to work out bond order,pick the same bond ineach structure and average the bond orderfor that bond

Number of canonical structures

Resonance in benzene.

C

C C

C

CCH

H

H H

H

H

C

C C

C

CCH

H

H H

H

H

There are two canonical structuresfor benzene, whichmeans that the C to Cbonds have a bond order of (2+1)/2 = 1.5.The benzene ring hasa very high stabilitydue to this resonance,which is calledaromaticity.

or

Short-hand versions for the benzene ring

8.7. Exceptions to the octet rule.

BF3. This can be written as F2B=F with three resonance structures. To complete its octet, BF3 readily reacts with e.g. H2O to form BF3.H2O. The actual structure of BF3 appears not to involve a double bond and does not obey the octet rule:

Possible resonancestructure for BF3, but is not importantas this wouldinvolve the very electronegativeF donating e’s to B

Best repre-sentation ofBF3 with B having only 6 electrons in its valence shell

Exceptions to the octet rule: free radicals

There are some molecules that do not obey the octet rule because they have an odd number of electrons. Such molecules are very reactive, because they do not achieve an inert gas structure, and are known as free radicals. Examples of free radicals are chlorine dioxide, nitric oxide, nitrogen dioxide, and the

superoxide radical:

nitric oxide chlorine dioxide

odd electrons

Exceptions to the Octet rule: Heavier atoms (P, As, S, Se, Cl, Br, I) may attain more than an octet of

electrons:

Example: PF5.

In PF5, the P atom has ten electrons in its valence shell, which occurs commonly for heavier non-metal atoms:

F

F

F

F

F

P

PF5

P has10 valenceelectrons

leave off Felectrons notshared with P

Many phosphorus compounds do obey the octet rule:

PF3 and [PO4]3- :

three blue electrons arefrom charge on anion

Some compounds greatly exceed an octet of electrons:

IF7 XeF6

(both I and Xe have 14 valence e’s)(Think about [XeF8]2-)