1. Introduction to molecular structure – Part I Contents: atomic orbitals - generalities models used in developing the molecular structure molecular orbital theory application for hydrogen-ion molecule diatomic molecules examples of molecular orbital diagrams

2. Mathematical description of electrons Schrödinger equation which for one electron system has the form: The solution of this equation are themselves mathematical formulas, but not differential equations. They describe the electrons as waves can be represented graphically. These graphs are 3D pictures that show electron density and are called orbitals or electron clouds. e

3. Atomic orbitals Waves added togehter can occur in two ways: When waves interact, they don't reflect off each other-- they combine. If the amplitudes of the waves are both positive or both negative, the combined wave will have a larger amplitude. destructive interference constructive interference If the waves have opposite amplitudes, the resulting wave will have a smaller amplitude. first-wave second-wave sum of the waves

4. Types of atomic orbitals 1s orbital +- -- + + atomic orbitals: s, p, d, f,g For example, in a simple lowest-energy state of hydrogen atom, the electrons are most likely to be found within a sphere around the nucleus of an atom. In a higher energy state, the shapes become lobes and rings, due to the interaction of the quantum effects between the different atomic particles z xy + z xy + 2s orbital Each p orbital has a node. The orbitals lobe are labeled (+) and (-) which referes to the signs of the wave function . 2p orbitals

5. Again to Shrödinger equation.... Possible to solve for one electron system, but not possible for molecules containing two or more electrons ...need for approximation Principals models of molecular structure 1. Molecular orbital theory 2. Valence bond theory Formation of chemical bonding where  represents the many-electron wave-function

6. Valence Bond Theory Explains the structures of covalently bonded molecules  ‘ how’ bonding occurs Principles of VB Theory:  Bonds form from overlapping atomic orbitals and electron pairs are shared between two atoms  Lone pairs of electrons are localized on one atom

7. Molecular Orbital (MO) Theory was developed in 1920’s explains the distributions and energy of electrons in molecules useful for describing properties of compounds –Bond energies, electron cloud distribution and magnetic properties basic principles of MO Theory –Atomic orbitals combine to form molecular orbitals –Molecular orbitals have different energies depending on overlaping type : Bonding orbitals (lower energy than corresponding atomic orbitals AO) Nonbonding orbitals (same energy as corresponding AO) Antibonding orbitals (higher energy than corresponding AO) Note: Both theories are limited case, but Chapter 9 treat a more exact theory

8. Molecular orbital theory - hydrogen molecule–ion H 2 + 1s 1  bond atomic orbitals molecular orbitals and 2  antibond molecular orbitals Hydrogen atom has only 1 electron situated on the s orbital

9. An application: the hydrogen molecule-ion H 2 + the simplest system: two nuclei and one electron the equation which characterize this system is given by: e AB rArA rBrB R kinetic energy of the electron attraction of the electron e by A attraction of the electron e by B repulsion between A and B Coordinates used to specify the hamiltonian for the hydrogen molecule-ion

10. the set of solution lead to construction of the molecular potential energy the molecular potential energy curves vary with internuclear distance, R..... and the equation set of solution Bonding: the electronic density transfer into the internuclear region: lowering of energy Antibonding: the electronic density undergo to a small expansion around nuclei increasing the energy

11. Bonding and antibonding of atomic orbitals  molecular orbital formation Atomic orbitals of the H 2 + are added in the same way as the waves. The in-phase addition of two 1s orbitals will form a molecular orbital with electron density between two nuclei  bonding orbital. The out-of-phase addition of the same type of orbitals will produce a molecular orbital with no electron density between two nucleii, called node  antibonding orbital.

12. Bonding and antibonding of atomic orbitals for H 2 + - interpretation 2  antibonding orbital 1  bonding orbital. In a bonding orbital, the nuclei are attracted to the accumulation of electron density in the inter nuclear region. In an antibonding orbital, the nuclei are attracted to an accumulation of electron density outside the internuclear region

13. Mathematical system in bond formation Theoretically, molecular orbitals can be formed by overlapping the wave function of the atomic orbitals. This overlaping procedure is constructed by forming linear combination of atomic orbitals, LCAO. Adition of the atomic orbitals give the bonding MO:  A,  B denote molecular orbitals  A,  B wave function for the atomic orbital of atom having the nucleus in A, B respectively; c A, c B electrons coeficients which show what part of the electron was involved in the bond formation bonding antibonding e AB rArA rBrB R Atkins notations

14. Liniar combination of atomic orbitals LCAO For species with many atoms, LCAO will have the form: In principle, we should use an infinite basis set for a precise recreation of the molecular orbital, but in practice only a finite basis set is used. In Atkins is assume that the members of the basis set are real and each one is normalize to 1. The optimum value of the coefficients are found by applying the variation principle, which means to solve the secular equation: fomation of basis set See section 6.10 from Atkins. H rs – matrix element of the hamiltonian S rs – overlap matrix element

15. The energies for LCAO-MO: The energy expression are derived in Atkins, chapter 8.3 Positive integral which is attributed to the total Coloumbic potential energy arising from the interaction of electron clound around A with nucelus in B and B with nucelus in A, respectively Interaction of the overlap electronic density with nucleus A AB j’ AB k’ A BA B

16. The molecular orbital energy level diagram of H + 2 in LCAO approximation E-E- E+E+ 22 11 molecular orbitals atomic orbitals Energy E 1s AA BB Diagram asymmetry: an antibonding orbital is more antibonding than a bonding orbital is bonding The bonding orbital is of lower energy than its atomic orbitals and the antibonding orbital is of higher energy Cal. with LCAO: R=2.5a 0 E=170kJ/mol Exp: R=2.0a 0 ; E=255kJ/mol

17. How we will use this theory? 1.we used the atomic orbitals from the hydrogen molecule-ion to construct the wave –functions for many electrons atoms, by adding them one by one into the calculations 2. we will use the orbitals of the hydrogen molecule-ion to build the MO for diatomic molecules

18. Diatomic molecules – angular momentum the spatial symmetries of atomic orbitals and the number of each symmetry type are determined by the angular momentum of the electron. the angular momentum vector for molecule case will lie along the bond axis. The quantum number in this case is denoted by and it is analogous to the quantum number m in the atomic case. The possible values for are: 0, 1, 2, 3...  correspondingly, the allowed values of the angular momentum about the internuclear axis are 0, ±1 (h/2p), ±2(h/2p), etc., or in general, ±l(h/2p). Thus when is different from zero, each energy level is doubly degenerate corresponding to the two possible directions for the component along the bond axis. The molecular orbitals are labelled according to the values of the quantum number l. When = 0, they are called s orbitals; when = 1, p orbitals; when = 2, d orbitals, etc. the molecule as a whole rotates in space and the nuclei contribute to the total angular momentum of the system. The nuclei and the electrons of a diatomic molecule can rotate around both axes which are perpendicular to the bond axis

19. Overlaping of 2p orbitals 2p x, 2p y, 2p z : the direction of orbitals along the axes. 1. Overlaping of 2p x orbitals: Constructive interference from the 2p x orbitals Destructive interference for the 2p x orbitals

20. 2. Overlaping of 2p y /2p z orbitals Constructive interference from the 2p y orbitals Destructive interference from the 2p y orbitals

21. Basic principle of molecular orbital formation: -Linear combination of atomic orbitals that have the same symetry species, i.e.span the same irreducible representation, within the molecular point group. As was shown in chapter 5.16 ( see M.H. Lecture), only orbitals of the same symetry species may have a nonzero overlap (S  0) and hence contribute to the bonding Diatomic molecules

22. Criteria for selecting the correct bonding Group theory provides techniques for selecting the atomic orbitals that may contribute to bonding, but others types of arguments must be used to decide whether these orbitals do in fact contribute and to what extent  two important criteria: 1.Atomic orbitals must be neither too diffuse or too compact (their overlaping is too diffuse to have a high significance) 2.The energies of the orbitals should be similar.

23. Filling of Orbitals The valence eletctrons of the atomic orbitals are used to fill in the molecular orbitals. The rules of filling in the MO are the same as filling in the AO Diatomic molecules – periodic table where only s and p orbitals are important 1.Homonuclear molecules – the same atomic species:ex: Li 2, O 2, N 2... 2.Heteronuclear molecules – different atomic species: ex: CO, HF...

24. HOMO and LUMO are acronyms for Highest Occupied Molecular Orbitals and Lowest Unoccupied Molecular Orbitals respectively. The energy level difference of the two (HOMO-LUMO) can (sometimes) serve as a measure of the excitability of the molecule: the smaller the energy, the easier it will be excited. The HOMO and LUMO state

25. AO MO 2s  2s The MO diagram for Li 2 Energy 2p  * 2s  2p  2p  * 2p  * 2p Energy The MO diagram for O 2 2p  2p  2p  * 2p  * 2p AO MO 2s  2s  * 2s 2p 12 valence eletrons 2 valence electrons HOMO LUMO HOMO

26. MO for CO molecule C has 4 valence electrons O has 6 valence electrons Carbon AO MO 2s  2s  * 2s 2p  2p  2p  * 2p  * 2p 2p Oxygen AO HOMO LUMO The CO molecule has 10 valence eletrons

27. ..... To be continued by Staffan in two weeks