Structure �and �Bonding-- I

Deepshikha

Department of Chemistry

Introduction

Important points

Hybridization

One common approach is called “hybridization” . It’s a bit of a kludge, but it gets us where we need to go.

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• We’re going to take the s orbital and mix it together with a certain number of p orbitals (1, 2, or 3). The total number of orbitals (s + p) will give us the total number of hybrid orbitals.

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• All the hybrid orbitals will be identical, and will orient themselves the maximum distance apart (tetrahedral in the case of CH4).

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• Any p orbitals that aren’t part of the hybrid will be “left over” as unhybridized p orbitals.

Thanks for reading! James�

In chemistry,  hybridisation (or hybridization) is the concept of mixing atomic orbitals into new hybrid orbitals (with different energies, shapes, etc., than the component atomic orbitals) suitable for the pairing of electrons to form chemical bonds in valence bond theory. Hybrid orbitals are very useful in the explanation of molecular geometry and atomic bonding properties and are symmetrically disposed in space.

Hybridization

Carbon's ground state configuration is 1s² 2s² 2p² or more easily read:

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The carbon atom can use its two singly occupied p-type orbitals, to form two covalent bonds with two hydrogen atoms, yielding the singlet methylene CH₂, the simplest carbene.

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The carbon atom can also bond to four hydrogen atoms by an excitation (or promotion) of an electron from the doubly occupied 2s orbital to the empty 2p orbital, producing four singly occupied orbitals.

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The energy released by the formation of two additional bonds more than compensates for the excitation energy required, energetically favouring the formation of four C-H bonds.

In chemistry, hybridization is the concept of mixing atomic orbitals to form new hybrid orbitals suitable for the qualitative description of atomic bonding properties. Hybridized orbitals are very useful in the explanation of the shape of molecules. It is an integral part of valence bond theory. 

For example, examine the energy level diagram for carbon:

�Fig 1: Energy diagram of carbon�

According to this diagram, carbon should only be able to make two bonds. However, carbon can make four bonds and this is how:

A molecule is much more stable if has lesser energy. When bonds are formed, energy is released and the system becomes more stable. If carbon forms 4 bonds rather than 2, twice as much energy is released and so the resulting molecule becomes even more stable. However, before carbon can do this it actually has to gain some energy by promoting an electron from the 2s to one of the2p orbital.

The energy required for carbon to promote an electron from the 2s to the 2p orbital is not very significant. Keeping this in mind, the carbon atom promotes an electron from the 2s orbital to the 2p_z orbital. Now that we've got four unpaired electrons ready for bonding, another problem arises. It’s seen that in  methane all the carbon-hydrogen bonds are identical, but our electrons are in different kinds of orbitals namely 2s, 2p_x, 2p_y and 2p_z. We can’t get four identical bonds unless we have four identical orbital. So how can we come up with four identical orbitals for carbon that will form four identical bonds with four hydrogen molecules? This is where hybridization comes in.

The electrons rearrange themselves again in a process called hybridisation. This reorganises the electrons into four identical hybrid orbitals called sp³ hybrids (because they are made from one s orbital and three p orbitals).�

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sp³ hybrid orbitals look a like half of a p orbital which attached at the nucleus and they arrange themselves in space so that they are as far apart as possible. This forms a tetrahedron with the nucleus in the centre.

sp³ Hybridization

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• Now ethane, H₃C-CH₃ as a model for alkanes in general:
• Both C are sp³  hybridised.
• 6 C-H σ bonds are made by the interaction of C sp³ with H1s orbitals (see the red arrows)
• 1 C-C σ bond is made by the interaction of C sp³ with another C sp³ orbital (see the green arrow

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sp² Hybridization

Other carbon compounds and other molecules may be explained in a similar way. For example, ethene (C₂H₄) has a double bond between the carbons.

For this molecule, carbon sp² hybridises, because one π (pi) bond is required for the double bond between the carbons and only three σ bonds are formed per carbon atom.

In sp² hybridisation the 2s orbital is mixed with only two of the three available 2p orbitals, usually denoted 2p_x and 2p_y. The third 2p orbital (2pz) remains unhybridised.

forming a total of three sp² orbitals with one remaining p orbital. In ethylene (ethene) the two carbon atoms form a σ bond by overlapping one sp² orbital from each carbon atom. The π bond between the carbon atoms perpendicular to the molecular plane is formed by 2p–2p overlap. Each carbon atom forms covalent C–H bonds with two hydrogens by s–sp² overlap, all with 120° bond angles. The hydrogen–carbon bonds are all of equal strength and length, in agreement with experimental data.

spHybridization

The chemical bonding in compounds such as alkynes with triple bonds is explained by sp hybridisation. In this model, the 2s orbital is mixed with only one of the three p orbitals,

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resulting in two sp orbitals and two remaining p orbitals. The chemical bonding in acetylene (ethyne) (C₂H₂) consists of sp–sp overlap between the two carbon atoms forming a σ bond and two additional π bonds formed by p–p overlap. Each carbon also bonds to hydrogen in a σ s–sp overlap at 180° angles.

Sp-hybridization