Fundamentals of Organic Molecules and Semiconductors

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1 Fundamentals of Organic Molecules and Semiconductors

2 Molecule 2

3 Periodic Table of the Elements 3

4 Carbon Carbon is found in every living creature. Elemental carbon can be black (graphite), or hard and beautiful (diamonds). Building block of fossil fuels (gasoline and oil) Basis of organic chemistry Crystalline forms Diamond (metastable) A wide-gap semiconductor Graphite Parallel hexagonal sheets Fullerine Semiconducting Fullerene (C 60 ) Graphite Diamond -Covalent 4

Building block of fossil fuels (gasoline and oil) Basis of organic chemistry Crystalline

5 Bonding Covalent (C, Si) Electron sharing 2 or more electrons shared by 2 or more atoms Ionic (GaAs) Electron transfer Van der Waals (Bonding between covalent molecules) London forces (attraction between molecules with no permanent dipole moment) Dipolar attraction (polar molecules) Hydrogen bonding (covalently bound hydrogen + negative dipole) 5

molecules) London forces (attraction between molecules with no permanent dipole

6 Covalent Bonding C-O bond 340 kj/mol 1.43 Å C-C bond 360 kj/mol 1.54 Å C-H bond 430 kj/mol 1.11 Å C=C bond 600 kj/mol 1.33 Å C=O bond 690 kj/mol 1.21 Å Sharing of electrons to achieve stable electron configuration Small difference in electronegativity of elements Bond energy : kj/mol Directional bond; between specific atoms in a specific direction, normally along the line connecting the two atoms that share a pair of electrons. 6

21 Å Sharing of electrons to achieve stable electron configuration Small difference in electronegativity of

7 Valence Bond Theory and Hybridization 1) bond formation by overlapping orbitals: A description of covalent bond formation in terms of atomic orbital overlap is called the valence bond theory. It gives a localized electron model of bonding: 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. 2) hybridization of atomic orbitals: How do a carbon with a s-orbital and three p-orbitals combined with four hydrogen (s orbitals) form four bonds and all four bonds are found to be : In 1931, Linus Pauling proposed that the wave functions for the s and p atomic orbitals can be mathematically combined to form a new set of equivalent wave functions called hybrid orbitals. : The mathematical process of replacing pure atomic orbitals with reformulated atomic orbitals for bonded atoms is called hybridization. In a hybridization scheme, the number of hybrid orbitals equals to the total number of atomic orbitals that are combined. The symbols identify the numbers and kinds orbitals involved. 7

electrons is concentrated in the region of orbital overlap.

8 Atomic Orbitals 8

9 sp 3 Hybridization sp 3 signifies one s and three p orbitals are combined. Mixing one s orbital with three p orbitals yields four equivalent sp 3 hybrid orbitals. 9

10 sp 3 Hybridization The bonding in methane Each of the four C-H bonds results from head-on (s) overlap of a singly occupied carbon sp 3 hybrid orbital with a singly occupied hydrogen 1s orbital. Sigma bonds are formed by head-to-head overlap between the hydrogen s orbital and a singly occupied sp 3 hybrid orbital of carbon. 10

11 sp 3 Hybridization The bonding in Ammonia 11

12 sp 2 Hybridization The molecular geometry is trigonal planar with bond angle = 120. To explain its geometry, we can use the following rational. sp 2 signifies one s and two p orbitals are combined. 12

13 sp Hybridization Now consider BeCl 2 which has linear molecular geometry determined experimentally. In hybridization scheme that best describes this compound is that The combination of one s and one p orbital gives two sp hybrid orbitals oriented 180 apart. Two unhybridized p orbitals remain and are oriented at 90 angles to the sp hybrids. 13

In hybridization scheme that best describes this compound is that The combination of

14 sp 3 d Hybridization To described hybridization scheme to correspond to the 5- and 6- electron-group geometries of VSEPR theory, we need to go beyond s and p orbitals and traditionally this meant including d orbitals. We can achieve the five half-filled orbitals and trigonal-bipyramidal molecular geometry through the hybridization of one s, three p and one d orbitals of valence shell into five sp 3 d hybrid orbitals. 14

15 sp 3 d 2 Hybridization In the same way, we can achieve the six half-filled orbitals and octahedral geometry through the hybridization of one s, three p and two d orbitals of valence shell into six sp 3 d 2 hybrid orbitals. 15

through the hybridization of one s, three p and two d

16 Hybrid Orbitals 16

17 Hybrid Orbitals and Multiple Covalent Bonds Sigma (ρ) bonds are characterized by Head-to-head overlap. Cylindrical symmetry of electron density about the internuclear axis. Pi (π) bonds are characterized by Side-to-side overlap. Electron density above and below the internuclear axis. 17

Cylindrical symmetry of electron density about the internuclear axis.

18 Molecular Orbitals Molecular orbitals (MOs) are mathematical equations that describe the regions in a molecule where there is a high probability of finding electrons Molecular orbitals (MOs) are essentially combinations of atomic orbitals two types exist, bonding and antibonding orbitals Molecular orbitals (MOs) are built up in the same way as atomic orbitals The hydrogen molecule Antibonding MO = region of diminished electron density Bonding MO = enhanced region of electron density 18

types exist, bonding and antibonding orbitals Molecular orbitals (MOs) are built up in the same way as atomic orbitals The

19 Molecular Orbitals 19

20 Molecular Orbitals for the 2p Electrons 20

21 Carbon-Carbon Single Bonds The bonding in Ethane Single bonds are always σ bonds, because σ overlap is greater, resulting in a stronger bond and more energy lowering. 21

22 Carbon-Carbon Double Bonds A double bond is made up of one sigma bond and one pi bond 22

23 Carbon-Carbon Double Bonds 23

24 Carbon-Carbon Double Bonds 24

25 Bonding in Aliphatic hydrocarbons 25

26 Carbon-Carbon Triple Bonds 26

27 Conjugated System A chemically conjugated system is a system of atoms covalently bonded with alternating single and multiple (e.g. double) bonds (e.g., C=C-C=C-C) in a molecule of an organic compound. This system results in a general delocalization of the electrons across all of the adjacent parallel aligned p-orbitals of the atoms, which increases stability and thereby lowers the overall energy of the molecule. 27

28 Conjugated System Adjacent, overlapping p orbitals allows for......more resonance...more electron delocalization...lower electron energy...greater stability Compound ethene max ,000 Consequences of p orbital overlap Atoms with p orbitals must be planar Partial pi bond(s) Barrier to rotation 1,3-butadiene ,000 50,000 1,3,5-hexatriene ,000 1,3,5,7-octatetraene ,000 1,3,5,7,9-decapentaene ,000 1,3,5,7,9,11-dodecahexaene 28

29 Benzene and Oligo-acene This kind of structure gives rise to two important resonance hybrids and leads to the idea that all three double bonds are delocalized across all six carbon atoms 29

30 Benzene 30

31 Benzene : Electronic Structure 31

32 Energy Level and Band Gap 32

33 Band Gap for Organic Materials 33

34 Some Semiconductors α quartertiophene α -4T α-sexithiophene α -6T S S S S S S S S S S Tetracene Pentacene π-conjugated Molecules (oligomers) Transport π π* overlap C 60 34

Chapter 9. Chemical reactivity of molecules depends on the nature of the bonds between the atoms as well on its 3D structure

Chapter 9. Chemical reactivity of molecules depends on the nature of the bonds between the atoms as well on its 3D structure Chapter 9 Molecular Geometry & Bonding Theories I) Molecular Geometry (Shapes) Chemical reactivity of molecules depends on the nature of the bonds between the atoms as well on its 3D structure Molecular