Sigma and pi Orbitals - Revision history

Cmditradmin: /* Covalent bonding */

2009-12-22T00:14:43Z

Covalent bonding

← Older revision		Revision as of 17:14, 21 December 2009
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	In the excited state the electrons are increasingly destabilized as the atoms get closer, the surface is repulsive. If you have two atoms infinitely far apart the system is neither stabilized nor destabilized. If you bring atoms to a correct geometry the ground state is stabilized. If you bring them too close you get a destabilized state. The excited state is always increasingly destabilized with decreasing distance.		In the excited state the electrons are increasingly destabilized as the atoms get closer, the surface is repulsive. If you have two atoms infinitely far apart the system is neither stabilized nor destabilized. If you bring atoms to a correct geometry the ground state is stabilized. If you bring them too close you get a destabilized state. The excited state is always increasingly destabilized with decreasing distance.
			[[category:molecular orbitals]]
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Cmditradmin: /* Effect of field on π electrons */

2009-12-01T19:32:48Z

Effect of field on π electrons

← Older revision		Revision as of 12:32, 1 December 2009
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	[[Image:chargetransfer.JPG\|thumb\|~~300px~~\|]]		[[Image:chargetransfer.JPG\|thumb\|400px\|The charge-separated form of this mole will react more to an applied field.]]
	σ electrons have electron density between the nuclei so they are tightly bound and do not change position as much when exposed to an external field. π electrons have the electron density above and below of the plane of the nuclei. When you apply a field you change the distribution and induce a dipole.		σ electrons have electron density between the nuclei so they are tightly bound and do not change position as much when exposed to an external field. π electrons have the electron density above and below of the plane of the nuclei. When you apply a field you change the distribution and induce a dipole.

Cmditradmin at 22:20, 15 September 2009

2009-09-15T22:20:31Z

← Older revision		Revision as of 15:20, 15 September 2009
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	~~<br clear='all'>~~		See [http://www.chem1.com/acad/webtext/chembond/cb08.html An Introduction to molecular orbital theory]

	=== Orbital Overlap ===		=== Orbital Overlap ===

Nsylvain: /* Covalent bonding */

2009-08-27T00:50:29Z

Covalent bonding

← Older revision		Revision as of 17:50, 26 August 2009
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	In the excited state the electrons are increasingly destabilized as the atoms get closer, the surface is repulsive. If you have two atoms infinitely far apart the system is neither stabilized nor destabilized. If you bring atoms to a correct geometry the ground state is stabilized. If you bring them too close you get a destabilized state. The excited state is always increasingly destabilized with decreasing distance.		In the excited state the electrons are increasingly destabilized as the atoms get closer, the surface is repulsive. If you have two atoms infinitely far apart the system is neither stabilized nor destabilized. If you bring atoms to a correct geometry the ground state is stabilized. If you bring them too close you get a destabilized state. The excited state is always increasingly destabilized with decreasing distance.

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Cmditradmin: /* Effect of field on π electrons */

2009-08-25T22:17:36Z

Effect of field on π electrons

← Older revision		Revision as of 15:17, 25 August 2009
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	[[Image:chargetransfer.JPG\|thumb\|300px\|]]		[[Image:chargetransfer.JPG\|thumb\|300px\|]]
	&~~Sigma~~; electrons have electron density between the nuclei so they are tightly bound and do not change position as much when exposed to an external field. Π electrons have the electron density above and below of the plane of the nuclei. When you apply a field you change the distribution and induce a dipole.		σ electrons have electron density between the nuclei so they are tightly bound and do not change position as much when exposed to an external field. π electrons have the electron density above and below of the plane of the nuclei. When you apply a field you change the distribution and induce a dipole.

Cmditradmin: Σ and Π Orbitals moved to Sigma and pi Orbitals

2009-08-25T22:11:17Z

Σ and Π Orbitals moved to Sigma and pi Orbitals

← Older revision	Revision as of 15:11, 25 August 2009
(No difference)

Cmditradmin at 21:32, 24 August 2009

2009-08-24T21:32:52Z

Cmditradmin at 21:31, 24 August 2009

2009-08-24T21:31:36Z

← Older revision		Revision as of 14:31, 24 August 2009
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	=== Effect of field on pi electrons ===		=== Effect of field on π electrons ===

	[[Image:PI_orbit_anim.gif‎]]		[[Image:PI_orbit_anim.gif‎]]
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	[[Image:chargetransfer.JPG\|thumb\|300px\|]]		[[Image:chargetransfer.JPG\|thumb\|300px\|]]
	Sigma electrons have electron density between the nuclei so they are tightly bound and do not change position as much when exposed to an external field. Pi electrons have the electron density above and below of the plane of the nuclei. When you apply a field you change the distribution and induce a dipole.		Sigma electrons have electron density between the nuclei so they are tightly bound and do not change position as much when exposed to an external field. Π electrons have the electron density above and below of the plane of the nuclei. When you apply a field you change the distribution and induce a dipole.

Cmditradmin: /* p-p π Bonding */

2009-08-19T20:25:10Z

p-p π Bonding

← Older revision		Revision as of 13:25, 19 August 2009
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			=== Effect of field on pi electrons ===

			[[Image:PI_orbit_anim.gif‎]]


			[[Image:chargetransfer.JPG\|thumb\|300px\|]]
			Sigma electrons have electron density between the nuclei so they are tightly bound and do not change position as much when exposed to an external field. Pi electrons have the electron density above and below of the plane of the nuclei. When you apply a field you change the distribution and induce a dipole.



			For a typical molecule you go from a neutral ground state to a charge separated state by application of an electric field. Quantum -mechanically the electric field causes a mixing of these two states.

	=== Covalent bonding ===		=== Covalent bonding ===

Cmditradmin: /* p-p σ Bonding */

2009-06-10T20:10:50Z

p-p σ Bonding

← Older revision		Revision as of 13:10, 10 June 2009
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	=== p-p σ Bonding ===		=== p-p σ Bonding ===

	[[Image:P_orbital_overlap.png\|thumb\|300px\|~~Constructive and destructive p orbital combination~~]]		[[Image:P_orbital_overlap.png\|thumb\|300px\|P atomic orbitals combine to form $sigma; molecular orbitals]]
	Two p orbitals can combine contructively or destructively. In constructive situation you have a lower energy bonding orbital. If they align destructively you have the higher energy antibonding orbital.		Two p orbitals can combine contructively or destructively. In constructive situation you have a lower energy bonding orbital. If they align destructively you have the higher energy antibonding orbital.
	In a σ bond, if you look down the axis between the atoms bonded to each other, the orbital will appear cylindrically symmetric. In such cases the atoms at the end of the bonds can rotate without breaking the bond. This is the case from the bonds between the s orbitals in hydrogen. P orbitals can overlap in two different ways. If they overlap head-on a sigma bond is formed.<br clear='all'>		In a σ bond, if you look down the axis between the atoms bonded to each other, the orbital will appear cylindrically symmetric. In such cases the atoms at the end of the bonds can rotate without breaking the bond. This is the case from the bonds between the s orbitals in hydrogen. P orbitals can overlap in two different ways. If they overlap head-on a sigma bond is formed.<br clear='all'>

← Older revision		Revision as of 14:32, 24 August 2009
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	This is sometimes referred to as Linear Combination of Atomic Orbitals (LCAO). When there is a destructive combination of the two waves there is a node between the two atoms where there is a zero probability of finding the electron. The first case is known as a bonding orbital and the latter case is known as an anti-bonding orbital.		This is sometimes referred to as Linear Combination of Atomic Orbitals (LCAO). When there is a destructive combination of the two waves there is a node between the two atoms where there is a zero probability of finding the electron. The first case is known as a bonding orbital and the latter case is known as an anti-bonding orbital.

	[[Image:H2_antibonding_orbital.png\|thumb\|300px\|Electrons from two atoms combine to form a sigma bonding molecular orbital.]]		[[Image:H2_antibonding_orbital.png\|thumb\|300px\|Electrons from two atoms combine to form a σ bonding molecular orbital.]]
	Note that the electron density is higher between the nuclei in the bonding molecular orbital and the that the orbital is stabilized, that is lower in energy relative to the two isolated hydrogren atoms. The sign of both wavefunctions is positive. There is a decreased electron density between the atoms in the antibonding molecular orbital and the orbital is destabilized, that is higher in energy than the two isolated hydrogen atoms. In very simple molecular orbital theory we treat the destabilization energy of the antibonding orbital as identical to the stabilization energy of the bonding orbital but in fact it is slightly more destabilized than the bond orbital is stabilized.		Note that the electron density is higher between the nuclei in the bonding molecular orbital and the that the orbital is stabilized, that is lower in energy relative to the two isolated hydrogren atoms. The sign of both wavefunctions is positive. There is a decreased electron density between the atoms in the antibonding molecular orbital and the orbital is destabilized, that is higher in energy than the two isolated hydrogen atoms. In very simple molecular orbital theory we treat the destabilization energy of the antibonding orbital as identical to the stabilization energy of the bonding orbital but in fact it is slightly more destabilized than the bond orbital is stabilized.
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	[[Image:P_orbital_overlap.png\|thumb\|300px\|P atomic orbitals combine to form $sigma; molecular orbitals]]		[[Image:P_orbital_overlap.png\|thumb\|300px\|P atomic orbitals combine to form $sigma; molecular orbitals]]
	Two p orbitals can combine contructively or destructively. In constructive situation you have a lower energy bonding orbital. If they align destructively you have the higher energy antibonding orbital.		Two p orbitals can combine contructively or destructively. In constructive situation you have a lower energy bonding orbital. If they align destructively you have the higher energy antibonding orbital.
	In a σ bond, if you look down the axis between the atoms bonded to each other, the orbital will appear cylindrically symmetric. In such cases the atoms at the end of the bonds can rotate without breaking the bond. This is the case from the bonds between the s orbitals in hydrogen. P orbitals can overlap in two different ways. If they overlap head-on a sigma bond is formed.<br clear='all'>		In a σ bond, if you look down the axis between the atoms bonded to each other, the orbital will appear cylindrically symmetric. In such cases the atoms at the end of the bonds can rotate without breaking the bond. This is the case from the bonds between the s orbitals in hydrogen. P orbitals can overlap in two different ways. If they overlap head-on a σ bond is formed.<br clear='all'>

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	[[Image:chargetransfer.JPG\|thumb\|300px\|]]		[[Image:chargetransfer.JPG\|thumb\|300px\|]]
	Sigma electrons have electron density between the nuclei so they are tightly bound and do not change position as much when exposed to an external field. Π electrons have the electron density above and below of the plane of the nuclei. When you apply a field you change the distribution and induce a dipole.		Σ electrons have electron density between the nuclei so they are tightly bound and do not change position as much when exposed to an external field. Π electrons have the electron density above and below of the plane of the nuclei. When you apply a field you change the distribution and induce a dipole.


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	When two atoms with duets or octets that have not been satisfied get within close proximity, they can interact in such a way as to create a bond by "sharing" electrons. Sharing of electrons can achieve a stable electron configuration corresponding to a noble gas configuration. On the right side of the diagram the hydrogen atoms are far apart and do not interact. As they move closer the system is stabilized by the formation of a bond through overlap of their orbitals and the lowering of the kinetic energy of the system. If the atoms get too close there will be an extremely strong nuclear - nuclear repulsion that greatly destabilizes the system. The correct balance between these interactions is found at the equilibrium bond length.		When two atoms with duets or octets that have not been satisfied get within close proximity, they can interact in such a way as to create a bond by "sharing" electrons. Sharing of electrons can achieve a stable electron configuration corresponding to a noble gas configuration. On the right side of the diagram the hydrogen atoms are far apart and do not interact. As they move closer the system is stabilized by the formation of a bond through overlap of their orbitals and the lowering of the kinetic energy of the system. If the atoms get too close there will be an extremely strong nuclear - nuclear repulsion that greatly destabilizes the system. The correct balance between these interactions is found at the equilibrium bond length.

	How does the potential energy surface for a hydrogen-hydrogen sigma bond in the ground state (where the electrons are in the lower orbital) vs the excited states (ie where one electron is in an antibonding orbital)? These two states have a different energy dependence as a function of distance. A σ bonding combination initially becomes increasingly stabilized as you move from infinity until the point where the internuclear interaction (repulsion between protons) causes the energy to increase greatly. This point for hydrogen occurs at a distance of about .74 Å. The bond strength is 104 Kcal/mol, which is a very strong bond. The equilibrium constant for hydrogen radicals vs bond hydrogen would be 99999999.... favoring the bound state.		How does the potential energy surface for a hydrogen-hydrogen σ bond in the ground state (where the electrons are in the lower orbital) vs the excited states (ie where one electron is in an antibonding orbital)? These two states have a different energy dependence as a function of distance. A σ bonding combination initially becomes increasingly stabilized as you move from infinity until the point where the internuclear interaction (repulsion between protons) causes the energy to increase greatly. This point for hydrogen occurs at a distance of about .74 Å. The bond strength is 104 Kcal/mol, which is a very strong bond. The equilibrium constant for hydrogen radicals vs bond hydrogen would be 99999999.... favoring the bound state.

	In the excited state the electrons are increasingly destabilized as the atoms get closer, the surface is repulsive. If you have two atoms infinitely far apart the system is neither stabilized nor destabilized. If you bring atoms to a correct geometry the ground state is stabilized. If you bring them too close you get a destabilized state. The excited state is always increasingly destabilized with decreasing distance.		In the excited state the electrons are increasingly destabilized as the atoms get closer, the surface is repulsive. If you have two atoms infinitely far apart the system is neither stabilized nor destabilized. If you bring atoms to a correct geometry the ground state is stabilized. If you bring them too close you get a destabilized state. The excited state is always increasingly destabilized with decreasing distance.