Molecular orbital diatomics

Question 1

molecular orbitals (MOs) describe

Answer 1

where the electrons are found in molecule

Question 2

shapes of the orbitals

Answer 2

relate to the electron density

Question 3

if the electron density is high between two atomic nuclei, a stable bond forms –

Answer 3

a bonding MO

Question 4

If the electron density is low between two atomic nuclei,

Answer 4

the positively charged nuclei repel and a bond does not form antibonding MO

Question 5

As the atoms come closer together,

Answer 5

overlap of the AOs forms a MO

Question 6

If the two AOs have the correct phases

Answer 6

, a bonding MO forms and the overall energy decreases

Question 7

Consider the example of two s-orbitals on neighbouring atoms

Answer 7

• If we bring together two orbitals with the same phase, the result is a bonding sigma orbital

Question 8

if the AO becomes too close

Question 9

if the AO becomes too close

Answer 9

Electrons resides in the same space, Pauli exclusion rule.

Question 10

if the AO becomes too close

Answer 10

Electrons resides in the same space, Pauli exclusion rule.

Question 11

Bond order =

Answer 11

½  (number of electrons in bonding MOs – number of electrons in antibonding MOs)

Question 12

This formula recognizes that electrons in antibonding MOs cancel out the bonding effect of
electrons in bonding MOs.

Answer 12

A single bond requires a net number of 2 bonding electrons.
A double bond requires a net number of 4 bonding electrons.

Question 13

Bond order H2

Answer 13

For H2
the bond order is 1 and there is a single bond

Question 14

The MO diagram for He2

Answer 14

• Each He atom contributes two 1s electrons to the occupied MO
  The electronic configuration of He2
  is 1s2
  1s*

There is no net bonding

Question 15

The electronic configurations of the Li atoms

Answer 15

The 1s orbitals are too close to the Li atom nuclei to
overlap with each other in Li2
* We call these core electrons and do not include them
in the MO diagram
* Hence we focus on the valence electrons (in this case
the 2s electrons)
* If we are using AOs with radial nodes such as 2s, there
is no need to include the radial nodes in the diagrams

Question 16

How to tell if the molecular orbital is s or 

Answer 16

PI - NODES
SIGMA - NO NODES

Question 17

There are two pi MOs with the same energy

Answer 17

• One pi MO forms from overlap of px + px
  (pi x
  ) and another from overlap of py + py
  (pi y
  )
• The two MOs that result have the same energy (they are degenerate) but are at 90o
  to each other

Answer 18

• Dashed lines show which AOs combine to make each MO

Question 19

• The best overlap comes from

Answer 19

AOs of similar size

Question 20

Guiding rules for drawing the MO diagrams:

Answer 20

• Horizontal lines represent AOs or MOs, and energy increases up the page
• We focus on the valence electrons most involved in bonding (so the “core” 1s orbitals are often not shown#

Question 21

• Each pair of interacting AOs

Answer 21

produces both a bonding and an antibonding MO

Question 22

• Hence the number of MOs formed

Answer 22

is the same as the starting number of AOs

Question 23

• The strongest interactions are between

Answer 23

AOs of similar energy

Question 24

For N2 and lighter homonuclear diatomic molecules of the second period elements

Answer 24

, a picture allowing hybridization of the 2s and 2p orbitals gives a better representation of the bonding

Question 25

in B, C and N vs. O, F and Ne

Answer 25

the 2s and 2p orbitals are closer in energy

Question 26

In homonuclear diatomic molecules (e.g. H2, N2, O2),

Answer 26

the MOs are also labelled with a subscript g or u

Question 27

g

Answer 27

stands for gerade and means symmetric with respect to the centre of the molecule

Question 28

u

Answer 28

u stands for ungerade and means antisymmetric with respect to the centre of the molecule