Hybridization of orbitals, 11 (1)

A description of covalent bond formation in terms of atomic orbital overlap is called the valence-bond method

1. Hybridization of Orbitals Dr. K. Shahzad Baig Memorial University of Newfoundland ( MUN) Canada Petrucci , et al. 2011. General Chemistry: Principles and Modern Applications. Pearson Canada Inc., Toronto, Ontario. Tro , N.J. 2010. Principles of Chemistry. : a molecular approach. Pearson Education, Inc.

2. Lewis theory of chemical bonding, have shortcomings. For example, why metals conduct electricity or how a semiconductor works. Some cases require more sophisticated approaches. One approach involves the s , p , and d atomic orbitals, or mixed-orbital or hybrid orbitals . A second approach involves the creation of a set of orbitals that belongs to a molecule as a whole . Electrons are then assigned to these molecular orbitals Limitations of Lewis theory of chemical bonding

3. c) Where the curve is flat, at the bottom of the potential energy well, there is no tendency either to attract or to repel . T he molecule is most stable at this internuclear distance (74 pm). The potential energy reaches its lowest value (-436 kJ/ mol ) d) Where the interaction is repulsive, the potential energy increases as the internuclear distance, r, decreases. The potential energy is 0 when the two H atoms are infinitely separated. the potential energy decreases as the internuclear distance, r, decreases, where the interaction is attractive.

4. There are several approaches to understanding bonding. The Lewis theory can be applied to determine Lewis structure . VSEPR theory makes it possible to propose molecular shapes that are generally in good agreement with experimental results. However, neither method yields: quantitative information about bond energies and bond lengths, Lewis theory has problems with odd-electron species and Situations When it is not possible to represent a molecule through a single structure (resonance).

5. A description of covalent bond formation in terms of atomic orbital overlap is called the valence-bond method. The creation of a covalent bond in the valence-bond method is normally based on the overlap of half-filled orbitals , but sometimes such an overlap involves a filled orbital on one atom and an empty orbital on another. Core electrons and lone-pair valence electrons retain the same orbital locations as in the separated atoms, and the charge density of the bonding electrons is concentrated in the region of orbital overlap the Valence-Bond Method

6. Hybridization of Atomic Orbitals The valence bond theory  to describe bonding in molecules in this bonds are considered to form due to the overlapping of two atomic orbitals on different atoms , each orbital containing a single electron. The localized  valence bonding theory  uses a process called hybridization. In hybridization atomic orbitals that are similar in energy, but not equivalent are combined mathematically to produce sets of equivalent orbitals that are properly oriented to form bonds. These new combinations are called hybrid atomic orbitals because they are produced by combining ( hybridizing ) two or more atomic orbitals from the same atom.

7. Oxygen has the electron configuration 1 s 2 2 s 2 2 p 4 , with two unpaired electrons (one in each of the two  2p  orbitals). Valence bond theory would predict that the 2 O–H bonds will form from the overlap of these two 2 p orbitals with the 1 s  orbitals of the hydrogen atoms. If this were the case, the bond angle would be 90°, because  p  orbitals are perpendicular to each other. For example, let us consider the water molecule, in which we have one oxygen atom bonding to two hydrogen atoms. Experimental evidence shows that the bond angle is 104.5°, not 90°. The prediction of the valence bond theory model does not match the real-world observations of a water molecule; a different model is needed.

8. Quantum-mechanical calculations suggest why the observed bond angles in H 2 O differ from those predicted by the overlap of the 1 s  orbital of the hydrogen atoms with the 2 p  orbitals of the oxygen atom. The mathematical expression known as the wave function,  ψ , contains information about each orbital and the wavelike properties of electrons in an isolated atom . When atoms are bound together in a molecule, the wave functions combine to produce new mathematical descriptions that have different shapes. This process of combining the wave functions for atomic orbitals is called hybridization. It is mathematically accomplished by the  linear combination of atomic orbitals , LCAO. The new orbitals that result are called  hybrid orbitals .

9. The following ideas are important in understanding hybridization : Hybrid orbitals do not exist in isolated atoms. They are formed only in covalently bonded atoms. Hybrid orbitals have shapes and orientations that are very different from those of the atomic orbitals in isolated atoms. A set of hybrid orbitals is generated by combining atomic orbitals. The number of hybrid orbitals in a set is equal to the number of atomic orbitals that were combined to produce the set. All orbitals in a set of hybrid orbitals are equivalent in shape and energy. The type of hybrid orbitals formed in a bonded atom depends on its electron-pair geometry as predicted by the VSEPR theory. Hybrid orbitals overlap to form σ bonds. Unhybridized orbitals overlap to form π bonds.

10. Bonding in H 2 O and NH 3 VSEPR theory describes a tetrahedral electron group geometry for four electron groups of H 2 O and NH 3 and suggest angles of o C. the experimentally observed bond angles of 104.5 o in H 2 O and 107 o in NH 3 SP 3 Hybridization

11. sp 2 Hybrid Orbitals Boron (group 13), has four orbitals but only three electrons in its valence shell. For most boron compounds, the appropriate hybridization scheme combines the 2s and two 2p orbitals into three sp 3 hybrid orbitals and leaves one p orbital unhybridized . The hybridization scheme corresponds to trigonal-planar electron group geometry and bond angles, as in BF 3

12. sp Hybrid Orbitals Beryllium, has 2 orbitals and only 2 electrons in its valence shell. In the hybridization scheme that best describes certain gaseous beryllium compounds, the 2s and one 2p orbital of Be are hybridized into 2 sp hybrid orbitals, and the remaining two orbitals are left unhybridized . The sp hybridization scheme corresponds to a linear electron-group geometry and a bond angle, 180 o as in BeCl 2 (g).

13. Sp 3 d and sp 3 d 2 Hybrid Orbitals To describe hybridization schemes for 5- and 6-electron group geometries of VSEPR theory, we need to go beyond the s and p subshells of the valence shell, and include d-orbital. PCl 5 sp 3 d SF 6 sp 3 d 5 The hybridization is not a real phenomenon, but an after-the-fact rationalization of an experimentally determined result.

14. The five regions of electron density around phosphorus in PCl 5  require five hybrid  sp 3 d  orbitals. These orbitals combine to form a trigonal bipyramidal structure with each large lobe of the hybrid orbital pointing at a vertex. There are also small lobes pointing in the opposite direction for each orbital (not shown for clarity).

15. To bond six fluorine atoms, the 3 s  orbital, the three 3 p  orbitals, and two of the 3 d  orbitals form six equivalent  sp 3 d 2  hybrid orbitals, each directed toward a different corner of an octahedron. Sulfur hexafluoride, SF 6 , has an octahedral structure that requires  sp 3 d 2  hybridization. The six  sp 3 d 2  orbitals form an octahedral structure around sulfur. Again, the minor lobe of each orbital is not shown for clarity.

16. https://chemistryonline.guru/hybridization-2/ Octahedral structure of sulfur hexafluoride

17. Hybrid Orbitals and the Valence-Shell Electron-Pair Repulsion (VSEPR) Theory To proceed to describe the bonding in molecules, the likely hybridization scheme for a central atom in a structure in the valence-bond method by writing a plausible Lewis structure for the species of interest selecting the hybridization scheme corresponding to the electron-group geometry using VSEPR theory to predict the probable electron-group geometry of the central atom

18. The hybridization scheme adopted for a central atom should be the one producing the same number of hybrid orbitals as there are valence-shell electron groups , and in the same geometric orientation. Thus, an hybridization scheme for the central atom predicts that four hybrid orbitals are distributed in a tetrahedral fashion. This results in molecular structures that are tetrahedral, trigonal-pyramidal, or angular , depending on how many hybrid orbitals are involved in orbital overlap and how many contain lone-pair electrons,