Transcript 14.1 Shapes of molecules and ions (HL)

14.1 Shapes of molecules
and ions (HL)
14.1.1
State and predict the
shape and bond angles using
the VSEPR theory for 5 and 6
negative charge centers.
Molecules with more than 4
electron pairs



Molecules with more
than 8 valence
electrons [expanded
valence shell]
Form when an atom
can ‘promote’ one of
more electron from a
doubly filled s- or porbital into an unfilled
low energy d-orbital
Only in period 3 or
higher because that is
where unused dorbitals begin
Why does this ‘promotion’ occur?
 When
atoms absorb energy (heat,
electricity, etc…)their electrons
become excited and move from a
lower energy level orbital to a
slightly higher one.
 How many new bonding sites formed
depends on how many valence
electrons are excited.
 Exceptions
to the octet rule. Shows
sulphur achieving 8, 10 and 12
valence electrons due to energy
input and excited electrons.
 http://www.saskschools.ca/curr_cont
ent/chem20/covmolec/exceptns.html
Trigonal Bipyramidal (5 pairs of V.E.)
Trigonal Bipyramidal
Normally would have 3 bp, but the lone
pair has moved from the p-orbital to
include the d-orbital, allowing for 2
additional bonding sites.
 Ex: PCl5

Octahedral (6 pairs of V.E.)
BrF5 is square pyramidal
SF6 is octahedral
XeF4 is square planar
Bond angles
 In
general, the greater the bond
angle, the weaker the repulsions.
 Equatorial- equatorial (120 o)
repulsions are weaker than axialequatorial (90o) repulsions.
– Equatorial: lie on the trigonal plane
(straight across)
– Axial: lies above and below the trigonal
plane (up and down)

Remember that lone
pairs cause more
repulsion than
bonding sites, so
expect the bond
angle to be changed
should there be lone
pairs, or double or
triple bonds involved
(multiple bonds also
cause more repulsion
than expected)
Practice:
1.
2.
3.
4.
5.
6.
7.
ClF3
PF5
XeO2F2
SOF4
SCl6
IF4+
ICl4-
1.
2.
3.
4.
5.
6.
7.
T-shaped
Trigonal
bipyramidal
Seesaw
Trigonal
bipyramidal
Octahedral
Seesaw
Square planar
14.2 Hybridization
.
14.2.1 Describe σ (sigma) and π (pi) bonds
14.2.2 State and explain the meaning of the
term hybridization
14.2.3 Discuss the relationships between
Lewis structures, molecular shapes and types
of hybridization (sp, sp2, sp3).
hybridization
the concept of mixing atomic orbitals to
form new hybrid orbitals
 Used to help explain some atomic bonding
properties and the shape of molecular
orbitals for molecules.
 The valence orbitals (outermost s and p
orbitals) are hybridised (mathematically
mixed) before bonding, converting some
of the dissimilar s and p orbitals into
identical hybrid spn orbitals
 We must know sp, sp2, and sp3 hydrid
orbitals

Hybrid orbitals
Carbon has 4
valence electrons.
 2 electrons paired
up in the s-orbital,
and 2 electrons
unpaired in the porbital.
 So why does it
commonly make 4
bonding sites?

 One
of carbon’s paired s-orbital
electrons is ‘promoted’ to the empty
p-orbital
 This produces a carbon in an excited
state which has 4 unpaired electrons
(4 equivalent bonding sites)
sp3 hybrid orbital


formed by mixing the
outermost s- and all
three outermost porbitals to form four
sp3 hybrids.
The furthest these
four [negatively
charged, and
therefore repulsive]
orbitals can get from
each other is the
corners of a
tetrahedron (109°).
Overlap four s-orbitals from four hydrogens (blue)
with four sp3 hybrids on carbon leads to formation
of bonds, each containing one electron from the
carbon and one from the hydrogen
Examples of sp3 hybrids


Methane, ammonia, water and hydrogen fluoride.
Note that the orbitals not involved in bonding to
hydrogen are still hybridised, but end up as lone
pairs of electrons (symbolised by the two dots in
the diagram above).
sp2 hybrid orbital
formed when only
one s- and two porbitals are
involved.
 This leaves one
remaining p
orbital, which may
be involved in
forming a double
bond.



The furthest these orbitals can get from one
another is a trigonal bipyramid, with the sp2
hybrids arranged at 120° to each other in a
plane.
This is characteristic of molecules with double
bonds.




Finally, sp hybrids are
formed using just one s
and one p orbital.
Two sp hybrids are formed
from them, and the two porbitals remaining may
contribute to a triple bond.
These arrange themselves
at the corners of an
octahedron, with the two
sp hybrids diametrically
opposite one another.
sp hybridisation is
characteristic of the triple
bond. (1 σ-bond and 2 π
(pi) bonds)
Sigma bond (σ-bond)
When s and/or hybrid orbitals overlap
'end-on', sigma bonds (σ) are formed
 They have a single area of electron
density between the nuclei of the two
atoms whose orbitals are overlapping.
 In the diagrams below, σ bond is shown

Sigma bond (σ-bond)
results from head-on overlap of orbitals
 electron density is symmetric about the
internuclear axis: between nuclei.

π (pi) bonds


p orbitals can overlap sideways too: when this
happens two lobes of electron density are formed
between the atoms.
From the diagram, you can see that the double
bond in ethene is composed of one σ plus one π
bond,
π (pi) bonds



results from sideways
overlap of orbitals
bonds resulting from
the combination of
parallel p orbitals
electron density is
above and below the
internuclear axis.
Predicting shape
 The
shape is dictated by the σ-bonds
and the non-bonding electron pairs
(lone pairs)
 π-bonds do not affect the shape of
the molecule (double bonds or triple
bonds)
– That’s why we refer to bonding sites
when using VSEPR, not paying attention
to whether it was single, double or triple
bonded.
14.3 Delocalization of
electrons
14.3.1 Describe the delocalization
of (pi) π- electrons and explain
how this can account for the
structure of some species
Delocalised electrons
The term 'delocalised' refers to an electron
which is not 'attached' to a particular atom
or to a specific bond.
 Delocalized electrons are contained within
an orbital that extends over several
adjacent atoms.
 Classically, delocalized electrons can be
found in double bonds and in aromatic
systems
 Double bonds = 1 sigma and 1 pi bond
 Delocalisation is often represented with
resonance structures or resonance hybrid

Resonance structures
the nitrate ion can be viewed as if it
resonates between the three different
structures above.
 Nitrate doesn’t change from one to the
next, but behaves as a combination of all
structures

Resonance is possible whenever a Lewis
structure has a multiple bond and an
adjacent atom with at least one lone pair.
 The following is the general form for
resonance in a structure of this type.

Practice
Try to show the
individual Lewis
structures for the
HCO3- ion
 Show its resonance
structure too

Practice drawing these resonance
structures:
1.
2.
3.
4.
5.
6.
NO3NO2CO32O3
RCOOBenzene (C6H6)


TOK
Kekule claimed that
the inspiration for the
cyclic structure of
benzene came from a
dream.
What role do the less
rational ways of
knowing play in the
acquistion of scientific
knowledge?
Bibliography and sites to visit

http://www.tutorvista.com/content/chemi
stry/chemistry-iii/chemical-bonding/typescovalent-bonds.php
– Good site on types of covalent bonds

http://www.mikeblaber.org/oldwine/chm1
045/notes/Geometry/VSEPR/Geom02.htm
– Used for expanded valence shell pictures

http://www.kentchemistry.com/links/bond
ing/lewisdotstruct.htm
– Puts the lewis diagrams together and explain
them. Including expanded shell

http://www.mpcfaculty.net/mark_bishop/r
esonance.htm
– Resonance structures pictures and notes

http://en.wikipedia.org/wiki/Delocalization
– Notes on delocalisation of electrons

http://www.steve.gb.com/science/atomic_
structure.html
– Amazing website for hybrid orbitals

http://library.thinkquest.org/C006669/dat
a/Chem/bonding/shapes.html
– Good review of all shapes