Lecture 11 Covalent Bonding Pt 3: Hybridization ( Ch. 9.5-9.13)

1. Lecture 11Covalent Bonding Pt 3: Hybridization (Ch. 9.5-9.13) Suggested HW: Ch 9: 25, 29, 39, 43, 72 (For 25 and 43, you are illustrating the hybridization of the atomic orbitals into hybrid orbitals and the overlapping of these hybrid orbitals as described in the examples provided)

2. Introduction • We now know that atoms can bond covalently through the sharing of electrons • VSEPR theory helps us predict molecular shapes. But, it does not explain what bonds are, how they form, or why they exist. • In ch 9, chemical bonding will be explained in terms of orbitals

3. Covalent Bonding Is Due to Orbital Overlap • In a covalent bond, electron density is concentrated between the nuclei. • Thus, we can imagine the valence orbitals of the atoms overlapping • The region of orbital overlap represents the covalent bond

4. Overlapping Valence Orbitals • Recall s and p orbitals (ch 5) S • S orbitals are spherical. L = 0, mL = 0 • Max of 2 electrons • P orbitals consist of two lobes of electron density. • L= 1, mL = -1, 0, 1 • (3 suborbitals) • Max of 6 electrons px py pz

5. Forming Sigma (σ) Bonds Energy σ + + + + H H H H Covalent bond stabilization (energy drop) 1s1 1s1 • Two overlapping atomic orbitals form a molecular bonding orbital. Plus sign indicates phase of electron wave, NOT CHARGE • Asigma(σ) bonding orbital forms when s-orbitals overlap. σ

6. Introduction to Hybridization • Imagine the molecule CH4. We know that carbon has 4 valence electrons (2s22p2). • However, when we fill our orbitals in order as according to Hund’s rule, we notice there are only enough unpaired electrons to make two bonds. Stay mindful of the fact that a covalent bond involves the sharing of unpaired electrons C X 2p2 4 H 2s2 ENERGY 1s1 1s1 1s1 1s1

7. sp3 Hybridization • So how does CH4form? How can carbon make 4 bonds? • To make four bonds, carbon hybridizes four of its atomic orbitals. This creates four equivalent sp3hybrid orbitals, each containing one unpaired electron. 2p2 ENERGY Four sp3hybrid orbitals 2s2 • The name “sp3” originates from the fact that the hybrid orbitals form as a result of the mixture of 1 s-orbital and 3 p-orbitals. Thus, each sp3orbital is 25% s character and 75% p character

8. Formation of Sigma Bonding Orbitals sp3hybrid orbitals C 1s1 1s1 1s1 1s1 ENERGY atomic s-orbitals 4H σbonding orbitals

9. Illustration of Orbital Hybridization pz s z + = z • The addition of an s-orbital to a pz orbital is shown above. The s orbital adds constructively to the (+)lobe of the pz orbital and adds destructivelyto the lobe that is in the opposite phase (-). The symbols indicate phase, not charge. • Whenever we mix a certain number of s and p atomic orbitals, we get the same number of molecular orbitals. This is called the principle of conservation of orbitals.

10. Illustration of sp3 Hybrid Orbitals and Orbital Overlap 4 σ-bonds The four hybrid orbitals arrange themselves tetrahedrally.

11. sp2 Hybridization • The BH3 molecule gives us an example of sp2 hybrid orbitals. • Once again, we have a situation where we don’t have enough bonding sites to accommodate all of the hydrogens. (Remember, B is electron deficient!) B 3 H X 2p1 ENERGY 2s2 1s1 1s1 1s1

12. sp2Hybridization • So, to make 3 bonding sites, 3 hybrid molecular orbitals are formed by mixing the 2s-orbital with two 2p-suborbitals. • This forms an sp2 orbital. Each of these three hybrid orbitals are one-thirds-character, and two-thirds p-character. unused 2p suborbital 2p1 ENERGY B Three sp2 hybrid orbitals 2s2

13. sp2 orbitals The result of adding one s and two p orbitals together is a trigonal planar arrangement of electron domains This figure illustrates the 3 hybrid orbitals combined with the unused 2p orbital, which is perpendicular to the hybrid orbitals.

14. sp2 Geometry and Bonding empty 2p orbital σ bond + + + H H H H B H H

15. sp Hybridization • Imagine BeH2 (the Be-H bond is covalent), with Be having the electron configuration: [He]2s2 • Here, we have a situation where no bonding electrons are available. To make two Be-H bonds, Be must create two hybrid orbitals by mixing two atomic orbitals (the 2s orbital and one of the 2p orbitals). This yields sp hybrid orbitals (50% s, 50% p) Be 2H 2p0 ENERGY X 2s2 1s1 1s1

16. sp Hybridization Be ENERGY unused 2p suborbitals 2p0 Two sphybrid orbitals 2s2

17. Hybridization of Lone Electron Pairs • Ex. What is the hybridization of Oxygen in H2O? • The valence electron configuration of O is [He]2s2 2p4 O 2p4 ENERGY 2H 2s2 1s1 1s1 As you see, there are two unpaired O electrons. Does this mean that these two p-suborbitals can overlap with the two Hydrogen 1s orbitals without hybridizing??

18. Hybridization of Lone Electron Pairs O •• •• • No!! The reason is that we now have two sets of lone pairs of electrons that are substantially different in energy (2s and 2p). The orbitals will hybridize to form degenerate(equal energy) sets of electrons. • Lone pair must always be equal in energy with each other, and with bonding electrons. 2p electrons 2s electrons O H H 2p4 ENERGY X 2H BAD!! 2s2 1s1 1s1

19. sp3 electrons Water has sp3 hybridization •• •• O O H H Lone pair Bonding electrons 2p4 Four sp3hybrid orbitals 2s2 2H ENERGY 1s1 1s1 H2O σ bonds

20. So What Do We Know So Far?

21. Double and Triple Bonding • How can orbital overlap be used to explain double and triple bonds? What kind of interactions are these? • Lets look at ethene, C2H4 The hybridization of each carbon is sp2because each is surrounded by three electron domains. The geometry around each C is trigonal planar. sp2 sp2 H H C C H H

22. Forming Double Bonds C unhybridized p-electron 2p2 sp2 hybrid orbitals H H 2s2 C C • We can see that for each carbon atom, we need three sp2 orbitals and three unpaired electronsto make three sigma bonds. But how is the double bond formed? H H

23. Double Bonds formed by simultaneous σ and π interaction • • + + + + H H H H The remaining p-electrons form a π bond. This bond forms due to attraction between the parallel p-orbitals. The like-phase regions are drawn toward one another and overlap. All double bonds consist of 1 σ-bond and 1 π-bond

24. Triple Bonds formed by 1 σ-bond and 2 π-bonds.Ex. HCN sp sp • • H C N • Can you draw the orbital diagram for this molecule?

25. Examples • How many σ and π bonds are in each of the following molecules? Give the hybridization of each carbon. • CH3CH2CHCHCH3 • CH3CCCHCH2

26. sp3d and sp3d2 hybridization • Atoms like S, Se, I, Xe… etc. can exceed an octet because of sp3d and sp3d2 hybridization (combination of ns, np, and nd orbitals where n>3). • This results in either trigonalbipyramidal or octahedral skeletal geometry sp3d sp3d2

27. Exceeding an Octet. Example: SF6 Energy 3d0 sp3d2 hybrid orbitals 3p4 Fluorine lone pair 3s2 S sp3 6 F SF6

28. Exceeding an Octet. Example: SF6 unpaired electron overlap F x 6 S 3 lone pair

29. Look Familiar ???

30. Examples: • What is the hybridization of the central atom? • CO2 • H2CO • CH3CCH • IF5 • PCl5 • SeOF4