Sigma ( s ) and pi ( π ) bonding in C 2 H 4

1. Sigma (s) and pi (π) bonding in C2H4 FIGURE 11-14 General Chemistry: Chapter 11

2. Class examples • 4. The ethanoic acid (“vinegar”) moleculae and the methyl ethanoate molecule (an ester) shown on the next slide contain C=O double bonds. What is the hybridization of the C and O atoms in these double bonds? (Mention acetone, acetaldehyde, formaldehye?)

3. Esters The distinctive aroma and flavor of oranges are due in part to the ester octyl acetate, CH3CO2CH2(CH2)6CH3 General Chemistry: Chapter 26

4. Carbon-Carbon Triple Bonds • The C≡C triple bond is explained using an sp hybridization scheme for C. We imagine distributing the 4 valence electrons (again singly) over the 2s and three 2p orbitals. One s orbital and one p orbitalare combined to form two hybrid sp orbitals. Two p orbitals on C (each containing a single electron) can be used to form two pi bonds.

5. sand π bonding in C2H2 FIGURE 11-16 General Chemistry: Chapter 11

6. Other Molecules with Triple Bonds • Carbon monoxide: C≡O • Hydrogen cyanide: H-C≡N • Methyl cyanide: H3C-C≡N • Cyanoacetylene: H-C≡C-C≡N • Aside: The organic cyanides (or nitriles) are found wherever people are smoking tobacco. Many cyanoacetylenes are found in dusty interstellar clouds (so is ethanol!).

7. Hybridization Summary for C Atoms

8. Alkenes and Alkynes General Chemistry: Chapter 26

9. Hybridization – Class Examples • We will draw structures for a range of organic molecules and determine which hybridization scheme can be used to describe the bonding for each C atom. These molecules will include saturated hydrocarbons, unsaturated hydrocarbons, alcohols, carboxylic acids, amines, aromatic compounds…….

10. Molecular Orbitals and Wave Properties of Electrons • We’ve mentioned that atomic orbitals can combine constructively to form a bonding molecular orbital. In the simplest case two H atoms are joined using a bonding molecular orbital. The H2 molecule has lower potential energy (or, is more stable) than the two isolated H atoms. “Destructive” combinations of atomic orbitals are also possible.

11. Molecular Orbital Theory • Atomic orbitals are isolated on atoms. • Molecular orbitals span two or more atoms. • LCAO • Linear combination of atomic orbitals. Ψ1 = φ1 + φ2 Ψ2 = φ1 - φ2 General Chemistry: Chapter 11

12. Electron Density in Bonding and Antibonding Orbitals • Bonding orbitals – considerable electron density between the bonded atoms. • Non-bonding orbitals – very little electron density between the bonded atoms (energetically unfavourable result). • Bonding and antibonding sigma orbitals are represented on the next slide.

13. Formation of bonding and antibonding orbitals FIGURE 11-20 General Chemistry: Chapter 11

14. The interaction of two hydrogen atoms according to molecular theory FIGURE 11-21 General Chemistry: Chapter 11

15. Basic Ideas Concerning MOs • Number of MOs = Number of AOs. • Bonding (lower energy) and antibonding (higher energy) MOs formed from AOs. • e-fill the lowest energy MO first (aufbau process) • Maximum 2 e- per orbital (Pauli Exclusion Principle) • Degenerate orbitals fill singly before they pair up (Hund’s Rule). General Chemistry: Chapter 11

16. Molecular Orbitals – Learning Objectives • 1. Construct molecular orbital diagrams for diatomic molecules composed of elements from the first period elements (H and He) and the second period elements (Li, Be, B, C, N, O F and Ne). This includes species with +ve and -ve charges. (Eg. O2+ and CN-). • 2. Label MOs in the MO diagram and show their relative energies. Indicate whether MOs are bonding or anti-bonding.

17. Molecular Orbitals – Learning Objectives • 3. Use the molecular formula (for neutral molecules and diatomic ions) and charge to determine the total number of electrons that we must accommodate using the MO picture. • 4. Distribute all of the electrons among the available MOs – starting with the lowest energy MOs (sound familiar?).

18. Molecular Orbitals – Learning Objectives • 5. After counting the number of electrons in both bonding and anti-bonding orbitals determine the bond order. • 6. Use the MO diagram (and the number of electrons in the various molecular orbitals) to determine whether a molecule is diamagnetic or paramagnetic.

19. Molecular Orbitals – Learning Objectives • 7. Understand a surprising feature of molecular orbital theory. We can accommodate all of the valence electrons in various molecular orbitals for a diatomic species and end up with a bond order of zero!

20. No. e- in bonding MOs - No. e- in antibonding MOs Bond Order = 2 Bond Order Stable species have more electrons in bonding orbitals than antibonding. General Chemistry: Chapter 11

21. Molecular Orbitals – Nomenclature: • For the simplest atoms (H, He, Li, Be) only 1s and 2s orbitals are occupied in the ground electronic state. The overlap of two 1s orbitals can only produce a sigma (σ) bond. In the H2 molecule, for example, two 1s atomic orbitals can combine to form a σ1s bonding molecular orbitaland a σ1s* anti-bonding molecular orbital. When 2p orbitals come into play we can form both σ and π molecular orbitals.

22. Simplest Diatomics – MO Diagrams • MO diagrams are initially a bit confusing because they represent the formation of chemical bonds using both a “before picture” (showing the relative energies of the various atomic orbitals) and an “after picture”(showing the relative energies of the molecular orbitals). We’ll illustrate this with the molecules H2, He2, H2+ and He2+.

23. BO = (1-0)/2 = ½ BO = (e-bond - e-antibond)/2 H2+ BO = (2-0)/2 = 1 H2 BO = (2-1)/2 = ½ He2+ BO = (2-2)/2 = 0 He2 Diatomic Molecules of the First-Period FIGURE 11-22 • Molecular orbital diagrams for the diatomic molecules and ions of the first-period elements General Chemistry: Chapter 11

24. Class Examples • Draw molecular orbital diagrams for Li2 and Be2. Using the MO diagrams determine the bond order for both molecules and, as well, indicate from the MO diagrams whether the molecules are diamagnetic or paramagnetic.

25. Molecular Orbitals of the Second Period Elements • First period use only 1s orbitals. • Second period have 2s and 2p orbitals available. • p orbital overlap: • End-on overlap is best – sigma bond (σ). • Side-on overlap is good – pi bond (π). General Chemistry: Chapter 11

26. Molecules with 2nd Period Atoms • The simplest possible molecular orbital diagram that one could imagine for second row elements having 2p electrons is shown on the next slide. This slide would necessarily apply only to homonucleardiatomics. Note the “symmetrical disposition” of bonding and nonbonding orbitals.

27. Possible molecular orbital energy-level scheme for diatomic molecules of the second-period elements FIGURE 11-25 (PART A) General Chemistry: Chapter 11

28. MO Diagrams - Surprises • The MO diagram presented on the previous slide does not adequately explain all properties of diatomic molecules formed from second period elements. Overlap of 2p atomic orbitals produces six MOs whose order energy order can vary with atomic number of the bonded atoms. General Chemistry: Chapter 11

29. MO Diagrams – Surprises – C2: • Two possible MO diagrams are illustrated for the C2 molecule on the next slide. The presentation of MOs here is similar to that used in drawing orbital diagrams for atoms. By experiment we know that the C2 molecule (4 valence electrons contributed by each C atom for a total of 8) is diamagnetic. Which of the MO diagrams accounts for this diamagnetism?

30. Experiment shows C2 to be diamagnetic, supporting a modified energy-level diagram Expected MO Diagram for C2 Modified MO Diagram for C2 General Chemistry: Chapter 11