INTRA MOLECULAR FORCES
- Intra
molecular forces are the attraction of atoms within the molecule. It is a strong
bond.
· An intra molecular force is any force
that holds together the atoms making up a molecule or compound.
- Attraction
between positive and negative ions in a crystal of an ionic compound.
- Covalent
bonds in molecular substances.
- Metallic
bonds
Types of Bonds
There
are four basic ways in which chemical combination occurs:
(1)
Ionic or electrovalent bond
(2)
Covalent bond
(3)
Coordinate bond
(4)
Metallic bond
Ionic or Electrovalent
Bond.
Ionic or electrovalent bonds are
formed by transfer of valence electrons from one atom to another. This type of
bond unites two atoms one of which has excess electrons than the stable number
(2 or 8), and the other is short of electrons.
For Ex - Sodium chloride
is a typical compound formed in this way.
Here the sodium atom (2, 8, 1)
transfers its excess electron to chlorine atom (2, 8, 7), and thus both attain
a stable inert gas type electron configuration. Sodium acquires the electron
configuration of neon (2, 8) and becomes positively charged. Chlorine acquires
the electron configuration of argon (2, 8, 8) and becomes negatively charged.
These oppositely charged ions are held together by electrostatic force of
attraction. This type of bond is commonly found in inorganic compounds.
Ionic
or electrovalent compounds are non volatile, soluble in water and possess high
melting points. Their aqueous solutions conduct electric current.
The sodium atom loses the one
electron in its outer shell to the chlorine atom, which uses the electron to fill
its outer shell. When this occurs, the sodium atom is left with a +1 charge and
the chlorine atom a -1 charge. The ionic bond is formed as a result of the
attraction of the two oppositely-charged particles.
Covalent Bond.
Ø Covalent bonds are
formed by mutual sharing of electrons. This type of bond unites two atoms, both
of which are short of electrons. The two atoms contribute one electron each and
then share the resulting pair of electrons.
For
Ex - Hydrogen is the simplest compound
formed in this way.
H. + xH -----------H ~H or H-H
Here
the two electrons are shared and give to each hydrogen atom the configuration
of helium.
Ø This type of bond
is termed covalent bond and is indicated by a line. Covalent bonds are commonly
found in organic compounds.
Ø A covalent bond
between two atoms results from the overlap of an orbital of one atom with an
orbital of another atom. When two orbitals overlap they share the same region in
space and a new orbital, called a Molecular Orbital (MO), is formed.
Ø Covalent compounds
are volatile, generally insoluble in water, but soluble in organic solvents. They
possess low melting and boiling points. Their solutions do not conduct electric
current.
Ø A covalent bond is
formed only if: (1) The two combining orbitals are half-filled and (2) The bonding orbitals approach each
other in proper alignment needed for an effective overlap. The electrons in the
bonding orbitals have opposite spins.
Types Of Covalent
Bonds
There
are two main types of covalent bonds or molecular orbitals:
(1)
Sigma (α)bonds
(2)
Pi(π) bonds
Sigma Bond
A sigma bond is formed by the
linear or end-to-end overlap of orbitals. They may be
Obtained
a) By
the overlap of two s orbitals,
b)
By the overlap of Px and an s orbital,
c)
By the overlap of two p orbitals,
A
pi (π) bond is formed by parallel or side-to-side overlap of orbitals. For
example,
There are two important differences
between sigma and pi bonds.
(1)
Pi electrons are loosely held than a pair of electrons in a α bond. They are of
higher energy. As a result, π bonds are more easily broken and are more
reactive than α bonds.
(2)
Rotation of atoms is not possible around a π bond. If any attempt is made to do
so, the lobes of p orbitals will no longer be coplanar and will not overlap to
form the π bond. This restriction in rotation around a π bond is responsible
for Cis and Trans isomerism in alkenes.
Unlike an ionic bond, a covalent
bond holds together specific atoms. Covalent bonding can be single covalent,
double covalent, or triple covalent depending on the number of pairs of
electrons shared. Figure 7 shows the bonding that occurs in the methane
molecule, which consists of four single covalent bonds between one carbon atom
and four hydrogen atoms.
Carbon-Carbon
Single Bond
Carbon atom has the wonderful property of
uniting with other carbon atoms through covalent bonds. This serves to
construct the carbon structure of organic molecules.
For
example - The molecules of hydrocarbons, ethane and propane contain two and
three carbon atoms respectively linked by covalent bonds.
Carbon-Carbon
Double Bond
In some compounds, two of the valencies of a
carbon atom may be satisfied by union with the two valencies of another carbon
atom.
For
example - In ethylene the two carbons are united by two covalent bonds.
Carbon-Carbon
Triple Bond
Sometimes two adjacent carbons are linked
together by three covalent bonds.
For
example - acetylene
Coordinate Bond.
Coordinate bond is also formed by
mutual sharing of electrons but in this case the two electrons that are shared
come from the same atom. A coordinate bond unites two atoms, one of which has a
spare pair of electrons and the other is short of a pair of electrons. The first
atom (donor atom) contributes one pair (lone pair) of electrons and the second
atom (acceptor atom) accepts it.
After the formation of the bond, the
lone pair of electrons is held in common. The coordinate bond is represented by
an arrow, pointing away from the donor atom.
For example - the coordinate bond
is found in the boron hydride-ammonia complex.
When both shared electrons in a
covalent bond come from the same atom, the bond is called a coordinate covalent
bond. Coordinate covalent bond is a single bond similar in properties to a
covalent bond.
For
example- negatively-charged chlorate ion.
The ion consists of one chlorine atom
and three oxygen atoms with a net charge of -1, and is formed with two
coordinate covalent bonds and one covalent bond. The chlorine atom has
effectively gained an electron through the covalent bond, which causes the
overall negative charge.
Metallic Bond
Metallic bonds are bonds where the atoms achieve a
more stable configuration by sharing the electrons in their outer shell with
many other atoms.
Elements of similar but relatively
low electro negativities (i.e., metals) form metallic bonds. It is these bonds
that are operative in a typical metal and that are responsible for metallic
properties: reflectivity, conductivity, malleability, and strength. Metallic
bonds form between atoms that have fewer valence electrons than they have
valence orbitals.
It would contribute its valence
electron to a bond with one or the other of the pair. The sodium atoms are
therefore driven to aggregate, using their relatively few valence electrons to
bond together as many nuclei as possible via the Coulomb attraction. The
aggregation is very regular in nature, resulting in an ordered arrangement of
metal atoms called a lattice. As a result of aggregation, each sodium
atom uses all of its valence orbitals in forming partial bonds with a number of
neighbours.
The nuclei are represented as
points, and the electrons as waves. That the electrons are delocalized waves is
consistent with the conductivity of metals, and with their flexibility and
strength. Putting stress on metals moves the nuclei, but the waves adjust,
maintaining bonding. Because the electrons are waves, they are everywhere in
the crystal at once. Thus the electrons absorb and reemit photons of all
energies in the visible region of the spectrum, giving metals their highly
reflective appearance.
Polar and non polar
Bonds
When a covalent bond is formed
between two atoms with different electro negativities, the electrons involved
in the bond are not shared. The atom with higher electro negativity pulls the
bonding electrons closer to it.
In other words, the electron density
in the molecular orbital would be greater around the atom with higher electro negativity.
The result of this displacement of the molecular orbital toward the more
electronegative atom will be that the more electronegative atom will acquire a
small negative and the less electronegative atom will acquire a small positive
charge. This is called polar bond.
Let
us now consider the C-CI covalent bond. Because chlorine is more
electronegative than carbon, the electro n density in the molecular orbital
would be higher around chlorine atom than around the carbon atom. Thus the
chlorine atom acquires a small negative charge and the carbon atom acquires a
small positive charge. This may be indicated in the following way.
When a covalent bond is formed
between two atoms with same e1ectronegativities (C-C, H-H, and F-F), the
electrons involved in the bond are shared equally. Electron density in the
molecular orbital binding the two atoms together is same around each atom.
There is no positive and negative end. Such a bond is said to be a Non polar
Bond.
Polar and Non polar
Molecules
When the shared pair of electrons
which are forming the bond in a molecule are not shared equally, the resulting
molecule will have a positive end and a negative end. This type of bond is a
polar covalent bond. The molecules are called dipolar or polar molecules.
A molecule is said to be a polar
molecule if it satisfies the following two conditions:
(1)
The molecule must contain one or more polar bonds.
(2)
The polar bonds must be so directed that there are separate centres of positive
and negative charges in the molecule.
For example –
a)
Carbon
dioxide (CO2) has two polar carbon-oxygen bonds, but the molecule is
not polar. The carbon dioxide molecule is linear and the centre of the two
negative charges is at the same place as the centre of the positive charge -
the carbon atom (Fig. 5.35).
b)
Water
molecule has two polar oxygen-hydrogen bonds. The H-O--H bond angle is 104°.
The centre of the positive charge is between the hydrogen atoms (Fig. 5.36). As
the centre of positive charge and the centre of negative charge do not
coincide, water molecule is polar.
c)
Carbon
tetrachloride has four polar carbon-chlorine bonds, but the molecule is not
polar. The Carbon tetra chloride is tetrahedral (all Cl-C-Cl angles are
109°28'), and the centre of four negative charges is at the same place as the
centre of the positive charge of the carbon atom (Fig. 5.37).
Examples
Dipole
moment
Polar molecules behave like small
magnets and possess a dipole moment. The dipole moment (µ) is the
product of the magnitude of the charges and the distance between the charges.
The
measure of molecular polarity is a quantity called the dipole moment (µ). Dipole moment is defined as magnitude of charge (Q)
times distance (r) between the charges.
µ = (Q)(r) Q charge in coulombs
(C) r distance in meters (m)
The product of the strength of either of the charges in anelectric dipole and the
distance separating the two charges. Itis expressed in coulomb meters. Dipole moment is a vectorquantity;
its direction is defined as toward the positive
charge.
Many molecules have such dipole moments due to non-uniform distributions
of positive and negative charges on the various atoms. Such is the case with polar compounds like hydrogen fluoride (HF), where electron density is shared unequally between atoms.
A molecule
with a permanent dipole moment is called a polar molecule. A molecule is polarized when it carries an
induced dipole. The physical chemist Peter J. W. Debye was the first scientist to study molecular dipoles
extensively, and, as a consequence, dipole moments are measured in units named debye in his honor.
With respect to molecules, there are three types of
dipoles:
Permanent dipoles: These occur when two atoms in a molecule have
substantially different electro negativity: One atom attracts electrons more than another, becoming
more negative, while the other atom becomes more positive.
Instantaneous dipoles: These occur due to chance when electrons happen to be more concentrated in one place than another in
a molecule, creating a temporary dipole.
Induced dipoles: These can occur when one molecule with a permanent dipole
repels another molecule's electrons, "inducing" a dipole moment in
that molecule.
The
simplest example of a dipole is a water molecule. A molecule of water is polar
because of the unequal sharing of its electrons in a “bent” structure. Since
oxygen has a higher electro negativity than hydrogen, the side of the molecule
with the oxygen atom has a partial negative charge while the hydrogen, in the
center, has a partial positive charge. Because of this, the direction of the
dipole moment points towards the oxygen.
It is expressed in the units of Debye and written as D
(where 1 Debye = 1 x 10-18e.s.u cm). nA dipole moment is a vector
quantity and is therefore represented by a small arrow with a tail at the
positive center and head pointing towards a negative center. In the case of a
Water molecule, the Dipole moment is 1.85 D, whereas a molecule of hydrochloric
acid is 1.03 D and can be represented as:
