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Chemical bondingLearning objectives:Students should be able to• Explain the shape of, and bond angles in , molecules by using the
qualitative model of electron-pair repulsion( including lone pairs),using as simple examples: BF3( trigonal); CO2 (linear); CH4 (tetrahedral); NH3 (pyramidal); H2 O(non-linear);SF6 (octahedral)
• Predict shapes and bond angles of molecules and ions following the steps below:
• (i) dot and cross diagrams to show total valence electrons( including the charge),
• (ii) state bonded pair electrons and lone pair electrons around the central atom
• (iii) predict the shape and bond angle• Draw the shape of molecules and ions in three dimensional diagrams • Explain expansion of octet for those atoms or ions which do not obey
octet rule .......continued on next slide
Chemical bonding• Continued from last slide.......• Describe covalent bonding in terms of orbital overlap, giving σ and π bonds using ethane and ethene as examples• Explain the shape of and bond angle of ethane, ethene and
benzene molecules• Explain the terms bond energy, bond length and relationship
between them• Using the examples of ethane and ethene ,relate the bond
energy with the reactivities of covalent bonds
Skills to be developed ( TGC)
• Students should be able to cultivate, develop and sharpen the following skills:
• Acquiring knowledge with understanding• Problem solving (numerical as well as structural)• Handling Information...... Gathering, analysing, evaluating, predicting ,applying
and synthesising information• Communication......Written as well as spoken
HYBRIDISATION OF ORBITALS - REVISIONHYBRIDISATION OF ORBITALS - REVISION
The ground state electronic configuration of a carbon atom is 1s22s22p2
1 1s
22s
2p
On providing a bit of energy, one of the s electrons can be promoted into a p orbital. The excited state configuration is now 1s22s12p3
1 1s
22s
2p
The process is favourable because of the arrangement of electrons; four unpaired electrons experience less repulsion and so the arrangement is more stable
HYBRIDISATION OF ORBITALS - REVISIONHYBRIDISATION OF ORBITALS - REVISION
The four orbitals (one s and three p orbitals) combine or HYBRIDISE to give four new orbitals. All four new hybridised orbitals are equivalent.
2s22p2 2s12p3 4 x sp3
HYBRIDISE
sp3
HYBRIDISATION
HYBRIDISATION OF ORBITALS - REVISIONHYBRIDISATION OF ORBITALS - REVISION
Alternatively, if only three orbitals (one s and two p orbitals) combine or HYBRIDISE ; they give three new orbitals. All three new hybridised orbitals are equivalent, but one of the 2p orbitals that remains unhybridised remains unchanged.
2s22p2 2s12p3 3 x sp2 2p
HYBRIDISE
sp2
HYBRIDISATION
In ALKANES, the four sp3 hybridised orbitals repel each other into a tetrahedral arrangement.
In ALKENES, the three sp2 hybridised orbitals repel each other into a trigonal planar arrangement and the 2p orbital lies at right angles to them
STRUCTURE OF ALKENES - REVISIONSTRUCTURE OF ALKENES - REVISION
Covalent bonds are formed by overlap of orbitals.
An sp2 hybridised orbital from each carbon overlaps to form a single C-C bond.
The resulting bond is called a SIGMA (δ) bond.
STRUCTURE OF ALKENES - REVISIONSTRUCTURE OF ALKENES - REVISION
After the formation of sigma bond, the two unhybridised 2p orbitals come close to each other and also overlap sideways. This forms a second bond; called a PI (π) bond.
For maximum overlap and hence the strongest bond, the 2p orbitals are in line.
This gives rise to the planar arrangement around C=C bonds.
STRUCTURE OF ALKENES - REVISIONSTRUCTURE OF ALKENES - REVISION
two sp2 hybridised orbitals overlap to form a sigma bond between the two carbon atoms
ORBITAL OVERLAP IN ETHENE - REVIEWORBITAL OVERLAP IN ETHENE - REVIEW
Two unhybridised 2p orbitals overlap to form a pi bond between the two carbon atoms
s orbitals in hydrogen overlap with the sp2 orbitals in carbon to form C-H bonds
the resulting shape is planar with bond angles of 120º
Molecular shapes
ELECTRON PAIR REPULSION ELECTRON PAIR REPULSION THEORYTHEORY
“THE SHAPE ADOPTED BY A SIMPLE MOLECULE OR ION IS THAT WHICH KEEPS REPULSIVE FORCES TO A MINIMUM”
ELECTRON PAIR REPULSION THEORYELECTRON PAIR REPULSION THEORY
Molecules contain covalent bonds. As covalent bond consist of a pair of electrons, each bond will repel the other bonds. Al
Bonds are further apart so repulsive forces are less
Bonds are closer together so repulsive forces are greater
Al
All bonds are equally spaced out as far apart as possible
Bonds will therefore push each other as far apart as possible to reduce the repulsive forces.
Because the repulsions are equal, the bonds will also be equally spaced
ELECTRON PAIR REPULSION THEORYELECTRON PAIR REPULSION THEORY
Because of the equal repulsive forces between bond pairs, most simple molecules, have standard shapes with equal bond angles.
However, the presence of lone pairs on the central atom affects the angle between the bonds and thus affects the shape.
O
All bonds are equally spaced out as far apart as possible to give minimum repulsive forces
REGULAR SHAPESREGULAR SHAPES
Molecules, or ions, possessing ONLY BOND PAIRS of electrons fit into a set of standard shapes. All the bond pair-bond pair repulsions are equal.
All you need to do is to count up the number of bond pairs and chose one of the following examples...
C
2 LINEAR 180º BeCl2
3 TRIGONAL PLANAR 120º AlCl3
4 TETRAHEDRAL 109.5º CH4
5 TRIGONAL BIPYRAMIDAL 90º & 120º PCl5
6 OCTAHEDRAL 90º SF6
BOND BONDPAIRS SHAPE ANGLE(S) EXAMPLE
A covalent bond will repel another covalent bond
BERYLLIUM CHLORIDEBERYLLIUM CHLORIDE
Cl ClBe180°
BOND PAIRS 2
LONE PAIRS 0
BOND ANGLE...
SHAPE...
180°
LINEAR
ClBe Be ClCl
Beryllium - has two electrons to pair up
Chlorine - needs 1 electron for ‘octet’
Two covalent bonds are formed
Beryllium still has an incomplete shell
Al
ALUMINIUM CHLORIDEALUMINIUM CHLORIDE
Cl
AlCl
Cl
Cl
Aluminium - has three electrons to pair up
Chlorine - needs 1 electron to complete ‘octet’
Three covalent bonds are formed;
aluminium still has an incomplete outer shell.
Al
ALUMINIUM CHLORIDEALUMINIUM CHLORIDE
Cl
Cl
Al
120°Cl
Cl AlCl
Cl
Cl
BOND PAIRS 3
LONE PAIRS 0
BOND ANGLE...
SHAPE...
120°
TRIGONAL PLANAR
METHANEMETHANE
BOND PAIRS 4
LONE PAIRS 0
BOND ANGLE...
SHAPE...
109.5°
TETRAHEDRAL
C H CH
H
H
H
109.5°
H H
C
H
H
Carbon - has four electrons to pair up
Hydrogen - 1 electron to complete shell
Four covalent bonds are formed
C and H now have complete shells
PHOSPHORUS(V) FLUORIDEPHOSPHORUS(V) FLUORIDE
FP
P
F
F
F
F
F
Phosphorus - has five electrons to pair up
Fluorine - needs one electron to complete ‘octet’
Five covalent bonds are formed;
phosphorus can make use of d orbitals to expand its ‘octet’
PHOSPHORUS(V) FLUORIDEPHOSPHORUS(V) FLUORIDE
FP
P
F
F
F
F
F
BOND PAIRS 5
LONE PAIRS 0
BOND ANGLE...
SHAPE...
120° & 90°
TRIGONAL BIPYRAMIDAL
120°
F
F
P
F
F
F
90°
SULPHUR(VI) FLUORIDESULPHUR(VI) FLUORIDE
FS
S
F
F
F
F
F
F
Sulphur - has six electrons to pair up
Fluorine - needs one electron to complete ‘octet’
Six covalent bonds are formed;
sulphur can make use of d orbitals to expand its ‘octet’
SULPHUR(VI) FLUORIDESULPHUR(VI) FLUORIDE
FS
BOND PAIRS 6
LONE PAIRS 0
BOND ANGLE...
SHAPE...
90°
OCTAHEDRAL
S
F
F
F
F
F
F
F
F F
FS
F
F
90°
SULPHUR(VI) FLUORIDESULPHUR(VI) FLUORIDE
FS
BOND PAIRS 6
LONE PAIRS 0
BOND ANGLE...
SHAPE...
90°
OCTAHEDRAL
S
F
F
F
F
F
F
IRREGULAR SHAPESIRREGULAR SHAPES
If a molecule, or ion, has lone pairs on the central atom,
the shapes are slightly distorted away from the regular shapes.
This is because of the extra repulsion caused by the lone pairs.
BOND PAIR - BOND PAIR < LONE PAIR - BOND PAIR < LONE PAIR - LONE PAIR
OO O
As a result of the extra repulsion,
bond angles tend to be slightly less as the bonds are squeezed together or pushed closer together
AMMONIAAMMONIA
HN NH H
HBOND PAIRS 3
LONE PAIRS 1
TOTAL PAIRS 4
• Nitrogen has five electrons in its outer shell
• It cannot pair up all five - it is restricted to eight electrons in its outer
shell
• It pairs up only three of its five electrons
• 3 covalent bonds are formed and a pair of non-bonded electrons is
left
• As the total number of electron pairs is 4, the shape is BASED on a
tetrahedron but it is not tetrahedral
AMMONIAAMMONIA
ANGLE... 107°
SHAPE... PYRAMIDAL
HN NH H
HBOND PAIRS 3
LONE PAIRS 1
TOTAL PAIRS 4
H
H
N
H
H
H
N
H
107°H
H
N
H
• The shape is based on a tetrahedron but not all the repulsions are equal
• LP-BP REPULSIONS > BP-BP REPULSIONS
• The N-H bonds are pushed closer together
• Lone pairs are not included in the shape
WATERWATER
H OH
HBOND PAIRS 2
LONE PAIRS 2
TOTAL PAIRS 4
O
• Oxygen has six electrons in its outer shell
• It cannot pair up all six - it is restricted to eight electrons in its outer
shell
• It pairs up only two of its six electrons
• 2 covalent bonds are formed and 2 pairs of non-bonded electrons are
left
• As the total number of electron pairs is 4, the shape is BASED on a
tetrahedron but is not tetrahedral
ANGLE... 104.5°
SHAPE... BENT
H
O
H
H OH
HBOND PAIRS 2
LONE PAIRS 2
TOTAL PAIRS 4
O
H
O
H
104.5°H
O
H
• The shape is based on a tetrahedron but not all the repulsions are the
same
• LP-LP REPULSIONS > LP-BP REPULSIONS > BP-BP REPULSIONS
• The O-H bonds are pushed even closer together
• Lone pairs are not included in the shape
WATERWATER
XENON TETRAFLUORIDEXENON TETRAFLUORIDE
FXe Xe
F
F
F
F
• Xenon has eight electrons in its outer shell
• It pairs up four of its eight electrons
• 4 covalent bonds are formed and 2 pairs of non-bonded electrons are
left
• As the total number of electron pairs is 6, the shape is BASED on an
octahedron
BOND PAIRS 4
LONE PAIRS 2
TOTAL PAIRS 6
XENON TETRAFLUORIDEXENON TETRAFLUORIDE
FXe Xe
F
F
F
F
F
F F
FXe
F
F F
FXe
ANGLE... 90°
SHAPE... SQUARE PLANAR
BOND PAIRS 4
LONE PAIRS 2
TOTAL PAIRS 6
MOLECULES WITH DOUBLE BONDSMOLECULES WITH DOUBLE BONDS
C O C OO
Carbon - needs four electrons to complete its shell
Oxygen - needs two electron to complete its shell
The atoms share two electrons
each to form two double bonds
The shape of a compound with a double bond is calculated in the same way. A double bond repels other bonds as if it was single e.g. carbon dioxide
MOLECULES WITH DOUBLE BONDSMOLECULES WITH DOUBLE BONDS
C O C OO
DOUBLE BOND PAIRS 2
LONE PAIRS 0
BOND ANGLE...
SHAPE...
180°
LINEAR
O OC180°
Double bonds behave exactly as single bonds for repulsion purposes so the shape will be the same as a molecule with two single bonds and no lone pairs.
A double bond repels other bonds as if it was single e.g. carbon dioxide