Consider an example of the simplest hydrocarbon molecular Methane. According to experimental observations, Methane molecule has 4 identical C-H bonds with equal length and equal bond energy. All the four hydrogen atoms are arranged in a manner such that the four hydrogen atoms form corners of a regular tetrahedron.
Based on the valence theory, a covalent bond is formed between two atoms in a molecule when there is an overlapping of half-filled atomic orbitals containing unpaired electrons. In the case of the methane molecule, we first write down the electronic configuration of each atom - C and H. Each carbon atom has two unpaired electrons in the 2px and 2py orbitals. Based on the valence theory, only two hydrogen molecules could be paired to the two unpaired electrons of the carbon atom and there will be a formation of only 2 C-H bonds in the molecule.
This will lead to an incomplete octet in the 2nd orbital of the carbon molecule 2pz orbital is unfilled and so the molecule should be unstable. However, we see that actually the methane molecule is extremely stable in nature and has 4 C-H bonds and not two. Hybridization concept explains the formation of identical 4 C-H bonds and the tetrahedral shape of the molecule.
According to this concept, when a carbon atom reacts with a hydrogen atom, the electrons in the carbon atom initially go into an excited state as shown here:. Post excitation, hybridization can be imagined as the process where these 4 excited s and the p orbitals combine together to give a homogenous mixture and divide themselves into 4 identical orbitals having identical energy. These new orbitals have been termed as hybridized orbitals. Since there one s orbital and 3 p orbitals have combined to form the hybrid orbital, the hybridized orbitals are called as sp 3 orbitals.
The energy of these hybrid orbitals lie in between the energy levels of the s and the p orbitals as shown here:. Each sp 3 hybrid orbitals has one unpaired electron. Since these 4sp 3 orbitals are identical in terms of energy, there is a tendency amongst these electrons to repel each other. To minimize the repulsion between electrons, the sp 3 hybridized orbitals arrange themselves around the carbon nucleus in a tetrahedral arrangement. The resulting carbon atom is termed as sp 3 hybridized carbon atom.
Overlap of each of the 4sp 3 orbitals of the hybridized carbon atom with the s orbital of the hydrogen atoms leads to the formation of methane molecule. Sp 2 hybridization results in trigonal geometry. In aluminum trihydride, one 2s orbital and two 2p orbitals hybridize to form three sp 2 orbitals that align themselves in the trigonal planar structure. The three Al sp 2 orbitals bond with with 1s orbitals from the three hydrogens through sp 2 -s orbital overlap.
Similar hybridization occurs in each carbon of ethene. For each carbon, one 2s orbital and two 2p orbitals hybridize to form three sp 2 orbitals. These hybridized orbitals align themselves in the trigonal planar structure. For each carbon, two of these sp orbitals bond with two 1s hydrogen orbitals through s-sp orbital overlap. The remaining sp 2 orbitals on each carbon are bonded with each other, forming a bond between each carbon through sp 2 -sp 2 orbital overlap.
This leaves us with the two p orbitals on each carbon that have a single carbon in them. These orbitals form a? Because a double bond was created, the overall structure of the ethene compound is linear. However, the structure of each molecule in ethene, the two carbons, is still trigonal planar. This formation minimizes electron repulsion. Because only one p orbital was used, we are left with two unaltered 2p orbitals that the atom can use.
These p orbitals are at right angles to one another and to the line formed by the two sp orbitals. Figure 1: Notice how the energy of the electrons lowers when hybridized. These p orbitals come into play in compounds such as ethyne where they form two addition?
This only happens when two atoms, such as two carbons, both have two p orbitals that each contain an electron. An sp hybrid orbital results when an s orbital is combined with p orbital Figure 2. We will get two sp hybrid orbitals since we started with two orbitals s and p.
These hybridized orbitals result in higher electron density in the bonding region for a sigma bond toward the left of the atom and for another sigma bond toward the right. In addition, sp hybridization provides linear geometry with a bond angle of o. In magnesium hydride, the 3s orbital and one of the 3p orbitals from magnesium hybridize to form two sp orbitals. The two frontal lobes of the sp orbitals face away from each other forming a straight line leading to a linear structure.
These two sp orbitals bond with the two 1s orbitals of the two hydrogen atoms through sp-s orbital overlap. The hybridization in ethyne is similar to the hybridization in magnesium hydride. For each carbon, the 2s orbital hybridizes with one of the 2p orbitals to form two sp hybridized orbitals.
The frontal lobes of these orbitals face away from each other forming a straight line. The first bond consists of sp-sp orbital overlap between the two carbons. Another two bonds consist of s-sp orbital overlap between the sp hybridized orbitals of the carbons and the 1s orbitals of the hydrogens.
Then they can become a part of a larger chemical structure. The beryllium atom contains all paired electrons and so must also undergo hybridization. One of the 2s electrons is first promoted to the empty 2p x orbital see figure below.
Now the hybridization takes place only with the occupied orbitals and the result is a pair of sp hybrid orbitals. The two remaining p orbitals p y and p z do not hybridize and remain unoccupied see Figure 6 below. The geometry of the sp hybrid orbitals is linear, with the lobes of the orbitals pointing in opposite directions along one axis, arbitrarily defined as the x-axis see Figure 7. Each can bond with a 1s orbital from a hydrogen atom to form the linear BeH 2 molecule.
Figure 7. The process of sp hybridization is the mixing of an s orbital with a single p orbital the pxorbital by convention , to form a set of two sp hybrids. The two lobes of the sp hybrids point opposite one another to produce a linear molecule. Other molecules whose electron domain geometry is linear and for whom hybridization is necessary also form sp hybrid orbitals.
Examples include CO 2 and C 2 H 2 , which will be discussed in further detail later. First a paired 2s electron is promoted to the empty 2p y orbital see Figure 8. This is followed by hybridization of the three occupied orbitals to form a set of three sp 2 hybrids, leaving the 2p z orbital unhybridized see Figure 9. The geometry of the sp 2 hybrid orbitals is trigonal planar, with the lobes of the orbitals pointing towards the corners of a triangle see Figure 9.
Each can bond with a 2 p orbital from a fluorine atom to form the trigonal planar BF 3 molecule. The process of sp 2 hybridization is the mixing of an s orbital with a set of two p orbitals p x and p y to form a set of three sp 2 hybrid orbitals.
Each large lobe of the hybrid orbitals points to one corner of a planar triangle. Other molecules with a trigonal planar electron domain geometry form sp 2 hybrid orbitals. Ozone O 3 is an example of a molecule whose electron domain geometry is trigonal planar, though the presence of a lone pair on the central oxygen makes the molecular geometry bent.
The hybridization of the central O atom of ozone is sp 2. Only read the boron section. Skip to main content. Covalent Bonding. Search for:. Hybrid Orbitals Learning Objectives Define hybridization. Describe sp 3 hybridization and covalent bond formation. Do you recognize this plant? Figure 1. Orbital configuration for carbon atom. Figure 2. Promotion of carbon s electron to empty p orbital. Figure 3.
0コメント