FSc Notes Chemistry Part 1 Chapter 6 Chemical Bonding Lecture 8
Dipole Moment (LL):
Introduction:
Dipole moment is the measure of separation of charges in a molecule.
In other words, the degree of polarity of a molecule is expressed in terms of
dipole moment.
Definition: The product of magnitude of Between two changes
of equal magnitude with opposite sign is equal to dipole moment. Ie
Dipole moment= charge x distance
Ll= q x d.
Explanation:
The dipole moment is a vector quantity and is represented by
an arrow which is directed from +ve pole the molecule.
For example in HCL molecule, the bonding electron pair is
not shared equally b/w H and Cl atoms. The Cl atom, being more electronegative
pulls the electron pair closer to it. Thus CL gets a slight or partial negative
while H gets a partial positive charge. Thus HCL is a polar molecule and is
indicated as,
+ S---------- - S
H---------------Cl
Units of Dipole Moment:
The common unit of dipole moment is Debye ( D). The SI unit
of dipole moment is " coulomb x meter " ( ie Cm )
1D = 3.335 x 10-30 Cm
Or
1D = 1 x 10-18 esu x cm
A Debye can be defined as:
The magnitude of dipole moment ( ll) when the charge (q) is
1 x 10-10 esu ( electrostatic unit ) and distance (d) is IAO (ie 10-8cm ) is
known as one debye
Ie ID = 1 x 10-10 esu charge x 1x 10-8 cm distance
ID = 1 x 10-10 esu x cm
Factors upon Which the Dipole Moment of a Molecule Depends
Dipole moment depends upon the following factors,
- Polarity of molecule
- Magnitude of charge
- Geometry of molecule.
Measurement of Dipole Moment
The dipole moment of a molecule is measured with the help of
an electric condenser. These are two plates in the condenser. The plates are
made of conducting material.
The sample molecule ( whose D.M is to be measured ) are
placed b/w the plates and the plates are connected with the opposite poles of a
battery. Thus an electric field is established whose strength is equal to the
applied voltage (V) divided by the distanced b/w the plates. As the polar
molecule, has opposite charges of equal magnitude, therefore they cancel out
the effect of each other, and thus the not charge of the dipole ( polar
molecule ) is always Zero. Thus the polar molecules placed b/w the oppositely
charged plates, are attracted by none of the plates. However, the polar
molecules rotate in b/w the plates, in such a way that the –ve pole gets
towards. The +ve plate while the +ve pole of the dipole gets towards the –ve
plate. Thus during this rotation of the dipole by the electric field, the
strength of the electric field decreases because of doing work.
The decrease in the electric field strength is noted. Then
from this decrease in the electric field strength, the dipole moment of the
molecule is measured.
Application of Dipole Moment:
D.M is very much helpful in explaining the structures or
geometries of molecules. The value of dipole moment of a molecule illustrates
its correct geometry. It is clear from the following examples.
1. Water ( H2O ) :
As H2O is a tri-atomic molecule, so it can have two types of
geometry ie either " linear " or "Angular or Bent " If water has linear geometry, then the polarity or dipole
moment of both its O—H bond will cancel out each other being equal and opposite
and hence the net D.M of H2O should be OD Ie
If H2O has an angular geometry, then the dipole moments of
the two O—H bonds is more opposite, thus they don’t cancel out each other and
hence the D.M of H2O in this can should be greater than O.
When the D.M of water is measured out, practically, it comes
out to be 1.84 d. So it confirms that H2O cannot have linear geometry. It has
always a bent or angular geometry.
2.Carbon Dioxide ( CO2 ) AND Sulphur Dioxide ( SO2 )
I. CO2 : As CO2 is tri- atomic molecule, so it can have
either a linear or angular geometry. If CO2 has a linear geometry, then the
dipole moment of two C==O bonds will cancel out each other, being equal and
opposite and thus the not dipole moment of CO2 will be Zero Debye. On the other
hand, if CO2 has angular or bent structure, then the dipole moments of the two
C==O bonds are no more opposite, so they will not cancel out each other and
thus CO2 should have some dipole moment.
But practically, the dipole moment of CO2 comes out to be
Zero Debye, so Co2 always has linear geometry and can never have angular
geometry.
ii. SO2 : it is also a tri- atomic molecule so according to
the above explanation it can have a linear or angular geometry.
As practically the dipole moment of SO2 molecule comes out
to be 1.62D, so SO2 has always an angular or bent structure and can never have
a linear structure.
3.Boron Trifluoride ( BF3) and Ammonia ( NH3 )
The experimentally determined dipole moment of BF3 is Zero
Debye. It is because of the fact that the three B—F bonds of BF3 molecule, are
arranged symmetrically around the boron atom in the same plane. Thus the bond
moments of the three B—F bonds cancel out each other's effect ie The
experimentally determined dipole moment of Ammonia ( NH3 ) is 1.47D. it is
because of the pyramidal geometry of NH3.
In NH3 molecule, the three hydrogen atoms lie in one plane
symmetrically with N atoms at the apex of the regular pyramid. The dipole
moments of the three N—H bonds, on addition give a not dipole moment to NH3
molecule. In addition, three is a lone pair of electrons on N atom & the
–ve change of this lone pair remains un-chanced.
Bond Energy & Bond Length:
Bond Energy:
The amount of energy evolved, when one mole of a molecule is
from its neutral atoms is known as bond energy of that molecule
Or
The amount of energy absorbed during breaking of bonds in
one mole of a molecule to produce neutral atoms, is known as bond energy of
that molecule. Bond energy is measured in the unit of Joules per mole ( J/mole
) or Kilo joules per mole ( Kj/ mole ).
The bond energy of a molecule shows the strength of the
bonds of that molecule the higher the bond energy of molecule, stronger will be
its bonds & vice versa.
It is important to note that polar bond has greater bond
energy than non- polar bond. It is because of the fact that in case of polar
bond, the shared pair of electrons is more closer to the nucleus of more
electronegative element. Thus this shared pair of es is attracted strongly by
one the two nuclei and hence more energy is required to separated these
electrons.
Thus we can say that the bond energy also tell us about the
extent of polarity of a molecule. For example the bond energy of HF is greater
than that of HCL which shows that the H—F bond is more polar than H—CL bond.
Bond Length:
The distance b/w the nuclei of two covalently bonded atoms,
is known as bond length. The bond length b/w two atoms is often ( but not
always ) independent of the nature of molecule. For example in alkenes, the
C—C, bond length is 154 pm. The C—C bond length in diamond is also 154pm.
As the electro-negativity difference b/w the bonded atoms
increases, the bond length decreases and hence the bond energy increases ( ie
bond becomes stronger ). Thus it is clear that a polar covalent bond is
stronger than a non—polar covalent bond.
In case of polar covalent bond, the departure from
additively of bond length occurs. For example, in SIH4, the SI—F bond length is
found to be 154—159 pm. But the covalent radii of Si and F, the bond length of
Si—F bond comes out to be 117 + 64 = 181 IPM, Thus the calculated bond length (
181 ) of Si—F bond is greater than the actual bond length of Si—F bond just
because of the electro-negativity difference of Si & F ( ie due to polarity
). Thus polarity affects the bond length.
Similarly, hybridization scheme, also plays a virtual role
in the shortening of bond length. Due to the involvement of S orbital in
hybridization, the bond length decreases between S orbital is close to the nucleus
and hence it makes the atom more electronegative. Greater the S character in a
hybrid orbital, greater will be the electeronegativity of the hybridized atom,
and hence shorter will be the bond length, formed from that hybrid orbital with
some other atom and hence stronger will be the bond.
As in case of ethane, each carbon undergoes SP3
hybridization, where the S character is 25% hence C atoms have normal electro
negativity and the C—C bond has larger length ( ie 154 pM ) In case of ethane (
C2H4 ) each C undergoes SP2 hybridization, where the S character is 33.3 %
hence C atom of ethane are a little more electronegative than those of ethane
and hence the C—C bond length of ethene is shorter than that of entane than C—C
bond length of ethane is 133 pm In case of ethane C2H2 each C undergoes SP
hybridization, where S character is 50 % hence the C atoms of ethane are most
electronegative as compared to those of ethane & ethane. Thus the C—C bond
length of C2H2 is 120 pm.
Here multiple bond formation in case of ehtyne & ethane
also plays some part in C—C bond length shortening.
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