1. To offer more in-depth explanations of chemical bonding more sophisticated concepts and theories are required 14.1 and 14.2 Hybridization 1

2. Understandings Covalent bonds form between atoms as a result of the overlap of atomic orbitals. There are two types of overlap Sigma ( σ ) bonds are f ormed by the direct end on end overlap of atomic orbitals. This results in the electron density concentrated between the nuclei of the bonded atoms. Pi (π) bonds are formed by the sideways overlap of atomic orbitals. This results in the electron density above and below the plane of nuclei of the bonding atoms. A hybrid orbital results from the mixing of different types of atomic orbitals on the same atom to form the same number of new equivalent hybrid orbitals that can have the same mean energy as the contributing atomic orbitals. 2

3. Applications You should be able to Predict whether sigma or pi bonds are formed from the linear combination of atomic orbitals. Explain the formation of sp 3, sp 2 and sp hybrid orbitals in methane (CH 4 ), ethene (C 2 H 4 ), and ethyne (C 2 H 2 ). Identify and explain the relationships between Lewis structures, electron domain geometry, molecular geometry and the types of hybridization. 3

4. Hybridization is a more in-depth explanation of chemical bonding and brings together Lewis theory Valence-bond theory Quantum Mechanics Model Hybridization is a mathematical model which relates the type of bonding in a molecule to its symmetry. 4

5. Lewis Theory A covalent bond is formed when a pair of electrons is shared between two bonding atoms. Electron density is greatest between the nuclei of the two bonded atoms. 5

6. Quantum mechanics theory Electrons are located in atomic orbitals, a region in space around the nucleus where there is a high probability of finding an electron. The principle quantum number (n) determines the size and energy of the atomic orbitals in a shell (n=1,2,3...) The angular momentum quantum number determines the shape of an atomic orbital (s, p, d and f) Magnetic Quantum Number determines the different orientations of the atomic orbital in space (x, y, z) 6

7. S atomic orbital Carbon, C - 1s 2 2s 2 2p 2 1 possible shape Spherical As the density of the dots increases the probability of finding an electron in that location increases

8. p atomic orbital Carbon, C - 1s 2 2s 2 2p 2 Dumbbell shape Three possible orientations 8

9. Covalent bonds form between atoms as a result of the overlap of atomic orbitals (s and p orbitals). Explains the electron domain geometry Two type of overlap sigma ( σ ) bond Formed by the direct end on end overlap of atomic orbitals. This results in the electron density concentrated between the nuclei of the bonded atoms. pi (π) bond Formed by the sideways overlap of atomic orbitals. This results in the electron density above and below the plane of nuclei of the bonding atoms Valence Bond Theory 9

10. Valance bond theory - Assumptions An optimum amount of overlap of atomic orbitals gives rise to the most energetically stable covalent bond. 10

11. Hybridization & Hybrid Orbitals Hybridization is the mixing s and p atomic orbitals on a centrally bonded atom. Hybridization produces hybrid orbital's that have lower energy and are more energetically stable. Hybrid orbitals take the shape of the combined atomic orbital’s and determine electron domain geometry of the molecule. The three types of hybrid orbitals studied in IB Chemistry are sp 3 sp 2 sp 11

12. sp 3 hybrid orbitals Combination of one s and three p atomic orbitals. The s and p orbital's combine (mix) to form four sp 3 hybrid orbital’s. Each sp 3 hybrid orbital has an energy that lies somewhere between the energy of the s and p orbitals that were used in the hybridization. Tetrahedral electron domain geometry. 12

13. Methane, CH 4 Carbon, C 1s 2 2s 2 2p 2 Hydrogen, H 1s 1 The sp 3 hybrid orbitals on the carbon atom overlap with the s orbital on four hydrogen atoms to form four single sigma bonds. Electron domain geometry - tetrahedral Molecular geometry – tetrahedral 13

14. Hybridization involving double & triple bonds In double and triple bonds the bonding electrons create too much repulsion. The quantum mechanics model proposes that the two electrons in a double bond and the four electrons in a triple bond are located 90% of the time in two sausage shaped lobes above and below the nuclei of the bonding atoms. Called a pi (π) bond. two electrons are located 90% of the between the nuclei of the bonding atoms. Called a sigma ( σ ) bond. 14

15. sp 2 hybrid orbitals Combination of s and p orbitals to make three sp 2 orbitals. Direct overlap of the sp 2 orbitals that leads to a sigma bond The unhybridized p orbital can overlap sideways with neighboring p orbitals to form a π bond. Trigonal planar electron domain geometry 15

16. Ethene C 2 H 4 Carbon, C 1s 2 2s 2 2p 2 Hydrogen, H 1s 1 The three sp 2 hybrid orbitals on each carbon atom overlap with the s orbital on two hydrogen atoms and the sp 2 orbital on the neighboring carbon atom to form three sigma bonds. The un-hybridized 2p orbital on each carbon atom overlap sideways to form a pi bond containing 2 electrons. Electron domain geometry – trigonal planar Molecular geometry – trigonal planar Double bond = 1 σ and 1 π bond 16

17. Ethene, C 2 H 4 17

18. sp hybrid orbitals Combination of s and p atomic orbitals to produce two sp hybrid orbitals arranged in a linear geometry. Only one p orbital is "used” to form a sigma bond. The two "left-over" p orbitals on the central atom overlap sideways with the p orbital on a neighboring atom and form either a double bond or triple bond. Linear electron domain geometry 18

19. Ethyne, C 2 H 2 Carbon, C 1s 2 2s 2 2p 2 Hydrogen, H 1s 1 The two carbon sp hybrid orbitals overlap and form a sigma bond with the neighboring carbon atom and a hydrogen atom. The two 2p un-hybridized orbitals from both carbon atoms overlap above and below the nuclei of the two atoms to form two pi bonds. Triple bond = 1 σ and 2 σ bonds 19

20. Ethyne, C 2 H 2 20

21. Bond strength When the number of electrons in the bond increases there is a greater attraction between these electrons and protons in the nucleus increasing the strength of the bond. This means that triple bonds with 6 electrons are stronger than double bonds with 4 electrons which are stronger than a single bond with 2 electrons. Hybridization – The implications 21

22. Reactivity Reactivity is related to the bond strength The electrons in a pi bond are on average further from the nuclei of the bonding atoms than the electrons of the sigma bond. This means that Pi bond electrons are not attracted as strongly to the positive nucleus of the bonding atoms. This means that less energy is required to break them in a chemical reaction, making triple bonds more reactive than double bonds which are more reactive than single bonds. Hybridization – The implications 22

23. BondEnthalpy change / KJmol -1 Outline of reactivity C-C 1σ1σ 346346 KJ of energy is needed to break a σ bond in a chemical reaction C=C 1 σ 1 π 614614 – 346 = 268 KJ of energy is needed to break a double bond (the π bond is broken) in a chemical reaction CΞCΞ 1σ2 π1σ2 π 839839 – (2x346) = 146 KJ of energy is needed to break a triple bond (the two π bonds are broken) in a chemical reaction 23

24. Bond length When the number of electrons in the bond increases there is a greater attraction between these electrons and protons in the nucleus of the bonding atoms decreasing the bond length This means that triple bonds with 6 electrons are shorter than double bonds with 4 electrons which are shorter than a single bond with 2 electrons. Hybridization – The implications 24

25. Number of covalent bonds Hybridization means a carbon atom can form four covalent bonds when it is hybridized. Stability The total energy of the new hybrid orbital is lower. This makes hybrid orbitals more stable. Hybrid orbitals are singly filled. This makes them more stable because repulsion between pairs of electrons is minimized (Hunds rule) Hybridization – The implications 25

26. Predicting the type of hybridization around a central atom 1. Draw the Lewis structure for the molecule. 2. Deduce the type of hybridization from the electron domain geometry. Tetrahedral – sp 3 hybrid Trigonal planar – sp 2 hybrid linear – sp hybrid 26

27. Ammonia, NH 3 Nitrogen, N 1s 2 2s 2 2p 3 Hydrogen, H 1s 2 Three sp 3 hybrid orbitals overlap with the s orbital on three hydrogen atoms to form three single covalent sigma bonds. The paired sp 3 hybrid orbital forms the lone pair of electrons. Electron domain geometry - tetrahedral Molecular geometry – trigonal pyramid 27

28. Water, H 2 O Oxygen, O 1s 2 2s 2 2p 4 Hydrogen, H 1s 1 The two paired sp 3 hybrid orbitals on the oxygen atom form the two non bonding pairs of electrons. The other two sp 3 hybrid orbitals overlap with the s orbital on two hydrogen atoms to form two single covalent sigma bond. Electron domain geometry - tetrahedral Molecular geometry – trigonal pyramid 28 www.sciencep hoto.com Water molecule, artwork C014/0007

29. Benzene, C 6 H 6 Carbon, C 1s 2 2s 2 2p 2 Hydrogen, H 1s 1 sp 2 hybridized carbon atoms Each sp 2 hybridized carbon atom overlaps to forms three sigma bonds with the adjacent carbon atoms and one hydrogen atom. The un-hybridized 2p orbital on each carbon atom overlap sideways to form a pi bond. Pi electrons become delocalized over three adjacent carbon atoms 29

30. Graphite Each sp 2 hybridized carbon atom overlaps to forms three sigma bonds with the adjacent carbon atoms. The un-hybridized 2p orbital on each carbon atom overlap sideways forming a delocalized pi bond. The delocalized electrons give rise to the electrical conductivity of graphite. Electronic domain geometry – trigonal planar 30

31. Borontrifluoride, BF 3 Boron, B 1s 2 2s 2 2p 1 Fluorine, F 1s 2 2s 2 2p 5 Exception to octet rule One s and two p orbitals combine to make three sp 2 orbitals. The three sp 2 hybrid orbitals overlap with the singly filled 2p on each fluorine atom form three sigma bonds. Electron domain geometry – trigonal planar Molecular geometry – trigonal planar 31

32. C 60 Fullerenes Each sp 2 hybridized carbon atom overlap to forms three sigma bonds with the adjacent carbon atoms. The un-hybridized 2p electron on adjoining C atoms overlap to form delocalized pi electrons. The delocalized electrons give rise to the partial electrical conductivity of fullerenes. Electron domain geometry - tetrahedral 32