Covalent Bonding: Orbitals

1. Covalent Bonding: Orbitals

2. Hybridization – mixing of two or more atomic orbitals to form a new set of hybrid orbitals. • Mix at least 2 nonequivalent atomic orbitals (e.g. s and p). Hybrid orbitals have very different shape from original atomic orbitals. • Number of hybrid orbitals is equal to number of pure atomic orbitals used in the hybridization process. • Covalent bonds are formed by: • Overlap of hybrid orbitals with atomic orbitals • Overlap of hybrid orbitals with other hybrid orbitals

3. Methane: known to have four identical C-H bonds. • How is this possible if the valence orbitals of a carbon atom are 2s and 2p orbitals?

4. Carbon adopts a set of atomic orbitals other than its “native” 2s and 2p orbitals to bond to the hydrogen atoms in methane. • Combine the carbon 2s and 2p orbitals to form “hybrid” orbitals. • The four new orbitals are sp3orbitals because they are formed from one 2sand three 2porbitals (s1p3).

5. The hybridization of the carbon 2s and 2p orbitals can be represented by an orbital energy-level diagram.

6. sp3 Hybridization • Combination of ones and threeporbitals. • Whenever a set of equivalent tetrahedral atomic orbitals is required by an atom, the localized electron model assumes that the atom adopts a set of sp3orbitals; the atom becomes sp3 hybridized. Tetrahedral set of Four sp3 Orbitals

7. sp2 Hybridization Combination of ones and twop orbitals. Gives a trigonal planar arrangement of atomic orbitals. One p orbital is not used. Oriented perpendicular to the plane of the sp2 orbitals.

8. The hybridization of the s, px, and py atomic orbitals results in the formation of three sp2 orbitals centered in the xy plane. • The large lobes of the orbitals lie in the plane at angles of 120 degrees and point toward the corners of a triangle.

9. An orbital energy-level diagram for sp2 hybridization. • Note that one p orbital remains unchanged.

10. Ethylene (C2H4) is an important starting material in the manufacture of plastics. • The C2 H4 molecule has 12 valence electrons and the following Lewis structure. • The double bond acts as one effective pair, so in the ethylene molecule each carbon is surrounded by three effective pairs. • An “effective pair” of electrons is either a lone pair, a single bond, a double bond, or a triple bond. • In other words, an effective pair is an electron group used for the purposes of determining molecular geometry.

11. sp2 hybridization can be used to account for the bonds in ethylene. • The three sp2 orbitals on each carbon can be used to share electrons. • In each of these bonds, the electron pair is shared in an area centered on a line running between the atoms. • This type of covalent bond is called a sigma (σ) bond. • In ethylene, the σ bonds are formed using sp2 orbitals on each carbon and the 1s orbital on each hydrogen atom.

12. How can we explain the double bond between the carbon atoms? • In the σ bond the electron pair occupies the space between the carbon atoms. • The second bond must therefore result from sharing an electron pair in the space above and below the σ bond. • This type of bond can be formed using the 2p orbital perpendicular to the sp2 hybrid orbitals on each carbon atom.

13. The parallelp orbitals can share an electron pair, which occupies the space above and below a line joining the atoms, to form a pi (π) bond. • Note that σ bonds are formed from orbitals whose lobes point toward each other, but π bonds result from parallel orbitals.

14. Figure (a): The orbitals used to form the bonds in ethylene. • Figure (b): The Lewis structure of ethylene.

15. A double bond always consists of one sigma (σ) bond, where the electron pair is located directly between the atoms, and one pi (π) bond, where the shared pair occupies the space above and below the σ bond. • Important!! • Whenever an atom is surrounded by three effective pairs, a set of sp2hybrid orbitals is required.

16. sp Hybridization • Combination of ones and onep orbital. • Gives a linear arrangement of atomic orbitals. • Two p orbitals are not used. • Needed to form the  bonds.

17. When ones orbital and onep orbital are hybridized, a set of twosp orbitals oriented at 180 degrees results.

18. The orbital energy-level diagram for the formation of sp hybrid orbitals on carbon.

19. Consider carbon dioxide, which has the following Lewis structure: • In the CO2 molecule the carbon atom has two effective pairs that will be arranged at an angle of 180 degrees. • We need a pair of atomic orbitals oriented in opposite directions (sp orbitals).

20. Two effective pairs around an atom will always require sp hybridization of that atom. • The sphybrid orbitals are used to form the σ bonds between the carbon and oxygen atoms in CO2.

21. In CO2 each oxygen atom has three effective pairs around it, requiring a trigonal planar arrangement of the pairs. • A trigonal set of hybrid orbitals requires sp2 hybridization so each oxygen atom is sp2 hybridized. • One p orbital on the oxygen is unchanged and is used for the πbond with the carbon atom.

22. The sp orbitals on carbon form σ bonds with the sp2 orbitals on the two oxygen atoms. • The remaining sp2 orbitals on oxygen hold lone pairs. • The π bonds between the carbon atom and each oxygen atom are formed by the overlap of parallel 2p orbitals.

23. dsp3 Hybridization • Combination of oned, ones, and threep orbitals. • Gives a trigonal bipyramidal arrangement of five equivalent hybrid orbitals.

24. Consider phosphorus pentachloride (PCl5) whose Lewis structure is shown below. • Five electron pairs requires a trigonalbipyramidal arrangement, therefore, we need a trigonal bypyramidal set of orbitals on phosphorus. • A set of five effective pairs around a given atom always requires a trigonal bypyramidal arrangement, which in turn requires dsp3 hybridization of that atom.

25. The five P-Cl σ bonds are formed by sharing electrons between a dsp3 orbital on the phosphorus and an sp3 orbital on each chlorine. • The other sp3 orbitals on chlorine hold lone pairs.

26. d2sp3 Hybridization • Combination of twod, ones, and threep orbitals. • Gives an octahedral arrangement of six equivalent hybrid orbitals.

27. Consider sulfur hexafluoride (SF6) whose Lewis structure is shown below. • Sulfur has six pairs of electrons and requires an octahedral arrangement of pairs and an octahedral set of six hybrid orbitals, or d2sp3 hybridization. • Fluorine has four pairs and is assumed to be sp3 hybridized. • Six electron pairs around an atom are always arranged octahedrally and require d2sp3 hybridization of the atom.

28. The Localized Electron Model: A Summary • Localized Electron Model • Draw the Lewis structure(s). • Determine the arrangement of electron pairs using the VSEPR model. • Specify the hybridorbitals needed to accommodate the electron pairs.

29. The relationship of the number of effective pairs, their spatial arrangement, and the hybrid orbital set required.