Hybridisation
When bonds are formed between atoms in the formation of molecules, there is a change in the nature of the Atomic Orbitals in each atom. Effectively, new molecular orbitals are formed, and the electrons in these orbitals are those of the molecule as a whole. The alteration of the structure of the atomic orbitals is called hydridisation, as it involves combining a number of orbitals to create an equal number of new orbitals, where each of the new hybrid orbitals have properties which are an average of those of the orbitals from which they were created.A number of different types of hybrid orbitals are known for carbon in the organic compounds.
- Hybridisation sp
- Hybridisation sp2
- Hybridisation sp3
Hybridisation sp
The simplest alkyne, Ethyne (i.e. Acetylene), has a linear structure, where the two carbon atoms and the two hydrogens atoms that are attached to these carbon atom lie along a straight line. The carbon to carbon and carbon to hydrogen bonds are arranged as far apart in space as possible. Thus, these bonds are at 1800 to each other. The geometry of this structure cannot be explained using the shape of the Atomic Orbitals on the carbon atom.The Electronic Configuration of carbon in the ground state (i.e. the lowest energy state) is 1s(2) 2s(2) 2p(2). If energy is supplied to raise one of the 2s electrons to a higher energy level to fill the vacant 2p orbital, the electronic configuration of carbon in the excited state, is 1s(2) 2s(1) 2p(3). More specifically, the electronic configuration of the excited carbon atom is 1s(2) 2s(1) 2px(1) 2py(1) 2pz(1). If we leave two of the 2p sub-orbitals (e.g. the 2py(1) and 2pz(1) sub-orbitals) to form the second and third bonds of the carbon to carbon triple bond, we can rearrange the other two sub-orbitals (i.e. 2s(1) and 2px(1)) to form two equivalent hybrid orbitals. When these new orbitals are arranged as far apart in space as possible, the new orbitals are arranged in a plane and are 180 degrees apart.
The particular hybridisation of the orbitals of carbon atoms in ethene, are called sp1 hybrids orbitals, (i.e. they arises from the hybridising of one sigma orbital and one pi orbitals). The bonds between the hydrogen atoms and carbons, and the first bond (of the double bond) between the carbon atoms of ethene are formed by the end-on overlap of these sp1 hybrids orbitals to form sigma bonds, and accounts for the 180 degree bond angles observed in this compound. The second and third bonds (of the triple bond) between the carbon atoms in ethyne is formed by the side-on overlap of the original atomic unhybridised orbitals (i.e. the 2py(1) and 2pz(1) orbitals) and these are pi bonds.
Hybridisation sp2
The simplest alkene, Ethene, has a planar structure, where the two carbon atoms and four hydrogens atoms that are attached to these carbon atom lie in a plane. The carbon to carbon and carbon to hydrogen bonds are arranged as far apart in space as possible. Thus, these bonds are at 120 degrees to each other. The geometry of this structure cannot be explained using the shape of the Atomic Orbitals on the carbon atom.The Electronic Configuration of carbon in the ground state (i.e. the lowest energy state) is 1s(2) 2s(2) 2p(2). If energy is supplied to raise one of the 2s electrons to a higher energy level to fill the vacant 2p orbital, the electronic configuration of carbon in the excited state, is 1s(2) 2s(1) 2p(3). More specifically, the electronic configuration of the excited carbon atom is 1s(2) 2s(1) 2px(1) 2py(1) 2pz(1). If we leave one of the 2p sub-orbitals (e.g. the 2pz(1) sub-orbital) to form the second bond of the carbon to carbon double bond, we can rearrange the other three sub-orbitals (i.e. 2s(1) 2px(1) 2py(1)) to form three equivalent hybrid orbitals. When these new orbitals are arranged as far apart in space as possible, the new orbitals are arranged in a plane and are 120 degrees apart.
The particular hybridisation of the orbitals of carbon atoms in ethene, are called sp2 hybrids orbitals, (i.e. they arises from the hybridising of one sigma orbital and two pi orbitals). The bonds between the hydrogen atoms and carbons, and the first bond (of the double bond) between the carbon atoms of ethene are formed by the end-on overlap of these sp2 hybrids orbitals to form s bonds (sigma bonds), and accounts for the 120 bond angles observed in this compound. The second bond (of the double bond) between the carbon atoms in ethene is formed by the side-on overlap of the original atomic unhybridised orbitals (i.e. the 2pz(1) orbital) and this pi bond prevents free rotation in the carbon to carbon axis.
Hybridisation sp3
The simplest alkane, Methane, CH4, has a tetrahedral structure, where the four hydrogens that are attached to the central carbon atom are arranged symmetrically about the carbon atom, and arranged as far apart in space as possible. The geometry of this structure cannot be explained using the shape of the Atomic Orbitals on the carbon atom.The Electronic Configuration of carbon in the ground state (i.e. the lowest energy state) is 1s(2) 2s(2) 2p(2). If energy is supplied to raise one of the 2s electrons to a higher energy level to fill the vacant 2p orbital, the electronic configuration of carbon in the excited state, is 1s(2) 2s(1) 2p(3). In this state, the carbon atom would be expected to produce three bonds of one sort (i.e. from the three Pi orbitals) and one bond of another sort (i.e. from the single Sigma orbital). However, as the four bonds in methane are known to be equivalent, it is necessary to postulate the existence of mixed or hybrid orbitals.
More specifically, the electronic configuration of the excited carbon atom is 1s(2) 2s(1) 2px(1) 2py(1) 2pz(1). If the four sub-orbitals in the second shell are re-arranged to form four equivalent hybrid orbitals, the result is a new second shell with four new orbitals. These new orbitals are arranged as far apart in space as possible. Thus, the new orbitals point to the corners of a regular tetrahedron, and the angle between any two hybrid orbitals 109.5 degrees.
The particular hybridisation of the orbitals of carbon atom, which results in the formation of the hybrid orbitals on the carbon atom in the methane molecule, is called sp(3) (i.e. they arises from the hybridising of one Sigma orbital and three Pi orbitals).
1s(2) 2s(1) 2p(3). In this state, the carbon atom would be expected to produce three bonds of one sort (i.e. from the three p orbitals) and one bond of another sort (i.e. from the single s orbital). However, as the four bonds in methane are known to be equivalent, it is necessary to postulate the existence of mixed or hybrid orbitals. More specifically, the electronic configuration of the excited carbon atom is 1s(2) 2s(1) 2px(1) 2py(1) 2pz(1). If the four sub-orbitals in the second shell are re-arranged to form four equivalent hybrid orbitals, the result is a new second shell with four new orbitals. These new orbitals are arranged as far apart in space as possible. Thus, the new orbitals point to the corners of a regular tetrahedron, and the angle between any two hybrid orbitals 109.5 degrees. The particular hybridisation of the orbitals of carbon atom, which results in the formation of the hybrid orbitals on the carbon atom in the methane molecule, is called sp(3) (i.e. they arises from the hybridising of one s orbital and three p orbitals). #$K+Hybridisation sp2 The simplest alkene, Ethene, C2H4, has a planar structure, where the two carbon atoms and four hydrogens atoms that are attached to these carbon atom lie in a plane. The carbon to carbon and carbon to hydrogen bonds are arranged as far apart in space as possible. Thus, these bonds are at 120 to each other. The geometry of this structure cannot be explained using the shape of the atomic orbitals on the carbon atom. The electronic configuration of carbon in the ground state (i.e. the lowest energy state) is 1s(2) 2s(2) 2p(2). If energy is supplied to raise one of the 2s electrons to a higher energy level to fill the vacant 2p orbital, the electronic configuration of carbon in the excited state, is 1s(2) 2s(1) 2p(3). More specifically, the electronic configuration of the excited carbon atom is 1s(2) 2s(1) 2px(1) 2py(1) 2pz(1). If we leave one of the 2p sub-orbitals (e.g. the 2pz(1) sub-orbital) to form the second bond of the carbon to carbon double bond, we can rearrange the other three sub-orbitals (i.e. 2s(1) 2px(1) 2py(1)) to form three equivalent hybrid orbitals. When these new orbitals are arranged as far apart in space as possible, the new orbitals are arranged in a plane and are 120 degrees apart. The particular hybridisation of the orbitals of carbon atoms in ethene, are called sp2 hybrids orbitals, (i.e. they arises from the hybridising of one s orbital and two p orbitals). The bonds between the hydrogen atoms and carbons, and the first bond (of the double bond) between the carbon atoms of ethene are formed by the end-on overlap of these sp2 hybrids orbitals to form s bonds (sigma bonds), and accounts for the 120 bond angles observed in this compound. The second bond (of the double bond) between the carbon atoms in ethene is formed by the side-on overlap of the original atomic unhybridised orbitals (i.e. the 2pz(1) orbital) and this p bond (pi bond) prevents free rotation in the carbon to carbon axis. #$K+Hybridisation sp The simplest alkyne, Ethyne, C2H2, (i.e. Acetylene), has a linear structure, where the two carbon atoms and the two hydrogens atoms that are attached to these carbon atom lie along a straight line. The carbon to carbon and carbon to hydrogen bonds are arranged as far apart in space as possible. Thus, these bonds are at 1800 to each other. The geometry of this structure cannot be explained using the shape of the atomic orbitals on the carbon atom. The electronic configuration of carbon in the ground state (i.e. the lowest energy state) is 1s(2) 2s(2) 2p(2). If energy is supplied to raise one of the 2s electrons to a higher energy level to fill the vacant 2p orbital, the electronic configuration of carbon in the excited state, is 1s(2) 2s(1) 2p(3). More specifically, the electronic configuration of the excited carbon atom is 1s(2) 2s(1) 2px(1) 2py(1) 2pz(1). If we leave two of the 2p sub-orbitals (e.g. the 2py(1) and 2pz(1) sub-orbitals) to form the second and third bonds of the carbon to carbon triple bond, we can rearrange the other two sub-orbitals (i.e. 2s(1) and 2px(1)) to form two equivalent hybrid orbitals. When these new orbitals are arranged as far apart in space as possible, the new orbitals are arranged in a plane and are 180 degrees apart. The particular hybridisation of the orbitals of carbon atoms in ethene, are called sp1 hybrids orbitals, (i.e. they arises from the hybridising of one s orbital and one p orbitals). The bonds between the hydrogen atoms and carbons, and the first bond (of the double bond) between the carbon atoms of ethene are formed by the end-on overlap of these sp1 hybrids orbitals to form s bonds (sigma bonds), and accounts for the 1800 bond angles observed in this compound. The second and third bonds (of the triple bond) between the carbon atoms in ethyne is formed by the side-on overlap of the original atomic unhybridised orbitals (i.e. the 2py(1) and 2pz(1) orbitals) and these are p-bonds (pi bonds).
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