Syructure Of Organic Molecules

Structure of organic molecules Molecular structures of organic compounds depend upon the chemical bonds. Formation of bonds may take place in different ways depending upon the arrangement of electrons in the outermost shell of combining atoms. The different type of bonds may be classified into the following categories: (i) Ionic Bonding: This type of chemical bonding is formed as a result of complete transfer of one or more electrons from one atom to other so that both atoms acquire inert gas electronic configuration (i.e. 1s2 or ns2 np6 outermost electronic configurations). The atom that loses the electron becomes positively charged. For example in the formation of sodium chloride, sodium atom loses its electron to acquire a stable octet and chlorine accepts the electron to complete eight electrons. Thus, x Na x cl Na cl 2,8,1 2,8,7 2,8 2,8,8 Na

+

cl

–

Nacl (ii) Covalent Bonding: This kind of bond is formed by mutual sharing of electron pairs between the atoms of the same or different elements. It is essential for sharing that the two electrons must have opposite spins.

For example: the formation of chlorine, two chlorine atoms combine to form a molecule by sharing a pair of electrons.

Formation of oxygen molecule.

Formation of nitrogen molecule.

Formation of water molecule.

The covalent bond is termed as single, double and triple bond when one, two and three pairs of electrons are shared respectively between two atoms. The number of electrons which an atom can share with others is known as the covalencyof the element. Thus hydrogen, oxygen and nitrogen have covalencies of 1, 2 and 3 respectively.

Orbital concept of covalent bond: Chemical bonding in organic compounds is concerned only with s and p orbitals. An orbital is that region in space where the electron charge density or probability of finding the electron is greatest by their geometric shapes. The S - atomic orbital is spherically symmetrical about the nucleus whereas the P atomic orbitals have a shape resembling a “dumb – bell” with two lobes disposed symmetrically about the nucleus along a line. This line is referred to as the orbital axis. The two lobes of the orbital have a position concentric with the atomic nucleus where the probability of finding an electron is zero. This point is called a node and the plane passing through it and perpendicular to the orbital axis is called nodal plane.

The three p – orbitals differ not only in shape from the s – orbital but also in directional properties. They are oriented perpendicular to each other along the three axes and therefore, are designated as px, py and pz orbitals.

Now according to orbital concept, a covalent bond is formed between two atoms when a half – filled valence orbital of one atom overlaps a half – filled valence orbital of another. It is essential that the two electrons which are involved in bond formation should have their spins in opposite directions. Thus a covalent bond is formed as a result of the progressive overlapping of the combining orbital. The atomic orbitals actually merge together to form a new orbital called molecular orbital. This orbital contains both the electrons which have now got paired. Types of covalent bonds: Two Types of covalent bonds are visualized as follow: (1) Sigma ( σ) bond (2) Pi ( π ) bond (1) Sigma

bond: A single bond formed between two atoms by end to end overlap or by linear overlapping of orbitals along their axes is known as sigma bond. Therefore sigma bond can be formed by linear overlapping of: (i) s – s orbitals (ii) s – p orbitals (iii) p – p orbitals, as shown in the diagram given below. s – s overlap

s – p overlap

p – p overlap

(2) Pi

( π ) bond :The bond formed between two atoms by the lateral or sidewise overlapping of atomic orbitals is known as a Pi bond. Therefore the Pi bond is formed by lateral overlapping of: (i)Py – Py orbitals (ii) Pz – Pz orbitals as shown in the diagram given below. Py – Py overlap

Pz – Pz overlap

Bond is weaker than a sigma bond. This is because of a relatively poor overlap of orbitals and thus the energy decrease in π bond formation is less than in sigma bond formation. Both sigma and Pi orbitals contain two electrons but the Pi electrons are easily attacked than the electrons in π

a sigma orbital. This is the reason why a double bond is more reactive than a single bond.

HYBRIDIZATION AND TETRA VALENCY OF CARBON 1. It is clear that the orbital theory does not offer a convincing explanation for the geometry of the molecules, since the formation of the compounds of carbon cannot be simply explained by the number of unpaired electrons in orbitals. The tetra valency of carbon in most of its compounds presents in intriguing problem and the attempt to describe the bonding in methane further brings up issues which are far more intriguing and cannot be solved easily. 2. Electron promotion approach: since the electronic configuration of carbon is 1s2 2s2 2p2, carbon has only two half – filled orbitals. Hence, according to the orbital theory carbon should be divalent and a compound having formula CH2 should be known. However, this is not so, it is a very reactive and unstable product. In actual practice carbon is always tetravalent and forms a very large number of stable compounds. In all those compounds, carbon has a covalency of four. To explain its tetravalency, it has been suggested that under the conditions of bond formation the electrons in the 2s orbitals of carbon atom are unpaired and one of them is promoted to the empty 2p orbital. This gives rise to the electronic configuration 1s2, 2p1, 2p3 as shown below.

3.

Therefore in the excited state, the carbon atom has four half – filled orbitals which would be available for the formation of four covalent bonds. Thus carbon atom would exhibit a covalency of four. Although the process of electron promotion justifies the tetra valency of carbon, the problem is still not solved because the four half – filled orbitals of carbon are not equivalent. A carbon atom has three p orbitals and one s orbital. It may be clearly explained by considering the formation of methane molecule. Here four hydrogen atoms form a single bond with the carbon. Now one of these bonds would involve the overlapping of 1s orbital of hydrogen with the 2s orbital of carbon to form one s – s bond while each of other three bonds would involve the overlapping of 1s orbital of hydrogen with a 2p orbital of carbon to form three s p bonds. But, in actual practice all the four bonds are found to be equivalent to one another in all respects and are directed towards the corners of a tetrahedron. Hybridization, to explain the equivalence of the four valencies of carbon, Pauling introduced the concept of hybridization. In this process, the four orbitals of carbon namely the 2s and 2px, 2py, 2pz combine together, losing their individuality and produce four new orbitals equivalent to each other in all respects. The new equivalent orbitals are called hybrid orbitals which have identical shape, size, energy and orientation is space. Orbit al

Mixing orbitals

Hybrid orbital produced

S

X

X

P

X

X

Types of Natur hybridizati e on Pure X Pure

X

Orientation Spherically Symmetrical Inclined at an

Sp Sp2

S+p S+p+p

Sp3

S+p+ p+ p

Two sp Three sp2

Four sp3

Hybri d Hybri d

Sp Sp2

hybri d

Sp3

angle of 90° Linear angle 180° Triangular plane, angle 120° Tetrahedral, angel 109° 28.

Thus by mixing s and p, two equivalent sp orbitals are obtained. By mixing s, p, p, three equivalent sp2 orbitals are obtained and by mixing s, p, p, p, four equivalent sp3 orbitals are obtained.

The characteristics of hybridized orbitals are as follows: 1- The number of hybrid of orbitals produced is equal to the number of atomic orbitals taking part in hybridization. 2- A hybrid orbital, like the atomic orbitals, cannot have more than two electrons in it, which two of opposite spins. 3- While distributing themselves in space, the hybrid orbitals assume the direction of the dominating orbitals. Since s orbital has no directional character of its own, the mixing p orbitals largely determine the shape and the direction of the resulting hybrid orbitals. Types of Hybridization Hybridization, the process of redistribution of energy in orbitals of an atom, is of following types: (a) Sp hybridization : The process of mixing of one s and one p – orbital to form two new hybrid orbitals is called sp – hybridization.

The characteristics of sp – hybrid orbitals are as follows: (i) The two sp – hybrid orbitals are linear i.e. lie in one plane. (ii) The angle between the two sp – hybrid orbitals is 180°. (iii) Each sp – hybridized orbital possesses 50 % s – character and 50% p – character. (iv) The 2py and 2pz orbitals remain perpendicular to the sp – hybridized orbitals.

(b) Sp

2

hybridization: The process of mixing of one s – orbital and two p – orbitals of an atom to form three equivalent hybrid orbitals directed towards the corners of an equilateral triangle is called sp2 hybridization.

The 2s, 2px and 2py orbitals hybridize before bonding to give three sp2- hybrid orbitals. The characteristics of sp2 – hybrid orbitals are as follows: (i) The three sp2 – hybrid orbitals lie one plane. (ii) The angle between any two sp2 – hybrid orbitals is 120°. (iii) All the three sp2 – hybrid orbitals are equivalent. Each sp2 – hybrid orbital possesses ⅓ s – character and ⅔ p – character. (iv) The remaining 2pz orbital which is left in its original state is oriented in a plane at right angles to the plane of the sp2orbitals.

(c)

Sp3 hybridization: The process of mixing of one 2s orbital and three 2p orbitals to form four new equivalent orbitals directed towards the corners of a tetrahedron is called sp3 hybridization. Thus in CH4, C2H6, C3H8, CCl4, CHCl3, CH3OH etc. hybridization of carbon is sp3.

The characteristics of sp3 – hybrid orbitals are as follows: (i) The four sp3 – hybrid orbitals of carbon are directed in space towards the corners of a tetrahedron. (ii) Each bond angle is 109° 28’. (iii) All the hybrid orbitals are equivalent. Each hybrid orbital possesses 25% s- character and 75% pcharacter. Now the different types of hybridization can be easily understood by following examples. The number of sigma bond on one carbon atom of a compound will indicate the type of hybridization. If a functional group is present in a compound then the total number of sigma bonds on the carbon atom of functional group or the carbon atom attached to functional group or a double or triple bond will indicate the type of hybridization. 1. For example in acetylene the number of sigma bonds on carbon atom attached to triple bond is two, hence there is sp hybridization. In the formation of acetylene molecule, the sp hybrid orbital of one carbon atom overlaps with that of the other to form a σ bond (sp – sp overlap). The other sp hybrid orbital of each carbon atom overlaps with s orbital of H atom to form another σ bond. The remaining two pure 2p orbitals (i.e. 2py and 2pz) along with one sp – orbital on each carbon atom are now available for side wise overlapping to form a π bond.

2. Formation of ethylene: H

H C

C H H Number of sigma bond between two carbon atoms = 1. Total number of sigma bonds in the molecule = 5. Number of sigma bonds on one carbon of ethylene = 3. .

. . Hybridization sp2 Shape triangular. 2 Here two sp hybridized carbon atoms form one sigma bond by head – on overlap. The remaining two sp2 orbitals of each carbon form σ bonds with H- atoms. The unhybridized 2pzorbitals of the two carbon atoms undergo a sidewise overlap forming a π bond.