INTRA MOLECULAR FORCES

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,

 Sigma bonds are symmetrical around the line drawn between the two nuclei. They are represented by a single line drawn between the atomic symbols. The electrons that occupy a sigma bond are called Sigma (α) Electrons.

 Pi-Bond.

A pi (π) bond is formed by parallel or side-to-side overlap of orbitals. For example,

             Like the p orbitals from which it is obtained, a π bond has two lobes. One half of the π bond lies above the plane containing the two nuclei and the other half lies below. The electrons that occupy π bond are known as the Pi (π) Electrons.

            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: