Note

If you cannot find what you are looking for. Please visit our sitemap

Wednesday, 21 October 2015

FSc Notes Chemistry Part 1 Chapter 6 Chemical Bonding Lecture 8

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,
  1. Polarity of molecule
  2. Magnitude of charge
  3. 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.

Written by: Asad Hussain

No comments:

Post a Comment