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Chemical Bonding 3 POLAR BONDS

Chemical Bonding 3 POLAR BONDS. University of Lincoln presentation. This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License. Definitions…. A HOMONUCLEAR BOND is a bond between two identical atoms

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Chemical Bonding 3 POLAR BONDS

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  1. Chemical Bonding 3 POLAR BONDS University of Lincoln presentation This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

  2. Definitions… • A HOMONUCLEAR BOND is a bond between two identical atoms • A HETERONUCLEAR BOND is a bond between different atoms This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

  3. Homonuclear & Heteronuclear bonds Hydrogen peroxide (H2O2) Hydrazine (N2H4) Ethane (C2H6) Homonuclear bonds Hetronuclear bonds This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

  4. Determining Bond Energies • Consider the 2 homonuclear diatomics H2and F2 • The bond energy of H–F would be expected to be the mean of the bond energies of H–H and F–F • Is this right? This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

  5. Bond Energies This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

  6. Anomalous Bond Energies This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

  7. Why are some heteronuclear bonds much stronger than expected? ? This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

  8. σ*(2s) Energy 2s 2s Li Li σ*(2s) SYMMETRICAL BONDS In a HOMONUCLEAR diatomic molecule, the electrons within the bond are shared equally between the two atoms – a symmetrical bond: The electrons sit in molecular orbitals which lie EQUI-DISTANT from each atom This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

  9. ASYMMETRICAL BONDS • In a HETERONUCLEAR diatomic molecule, the electrons within the bond are NOT always shared equally between the two atoms – an asymmetrical bond. • In an assymetrical bond, the electrons sit closer to one atom than the other, leading to a POLAR BOND: + – H–F The electrons are sitting closer to the F atom This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

  10. Why does this happen? This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

  11. Electronegativity Pauling defined ELECTRONEGATIVITY as: “the power of an atom in a molecule to attract electrons to itself” This is an atomic property, but only applies when the atoms are in a bond This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

  12. Electronegativity H 2.2 Li 1.0 Be 1.6 B 2.0 C 2.6 N 3.0 O 3.4 F 4.0 Na 0.9 Mg 1.3 Al(III) 1.6 Si 1.9 P 2.2 S 2.6 Cl 3.2 K 0.8 Ca 1.0 Ga(III) 1.8 Ge(IV) 2.0 As(III) 2.2 Se 2.6 Br 3.0 Rb 0.8 Sr 0.9 In(III) 1.8 Sn(IV) 2.0 Sb 2.1 Te 2.1 I 2.7 Tl(III) 2.0 Bi 2.0 Po 2.0 At 2.2 Cs 0.8 Ba 0.9 Pb(IV) 2.3 The higher the electronegativity, the stronger the ‘pulling’ power of the atom within a bond This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

  13. …When electrons are held tightly by an atom in a bond, due to the high electronegativity of that atom, the bond is much harder to break So, why are some heteronuclear bonds much stronger than expected? This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

  14. Examples of Polar Bonds - + The slight charges on each end of the molecule lead to electrostatic attraction between adjacent molecules – HYDROGEN BONDING + + – - + This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

  15. Definition… • A HYDROGEN BOND is an interaction between a hydrogen atom attached to an electronegative atom, and an electronegative atom which possesses a lone pair of electrons The strongest hydrogen bonds involve the first row elements F, O or N This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

  16. HYDROGEN BONDING ( ) H–F H–F H–F H–F H–F This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

  17. Hydrogen bonding affects the physical properties of molecules with polar bonds NH3, H2O and HF all have anomalously HIGH boiling points, since extra energy is needed to break the hydrogen bonds This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

  18. Can Molecular Orbital Theory account for polar bonds? ? This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

  19. A quick recap… ATOMIC Orbitals MOLECULAR Orbitals H + H H2 This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

  20. Electronic configuration of 9F is: 1s2 2s2 2p5 (9 electrons) The F atom needs 1 more electron to give it a full valence shell (8 outer electrons)– it does this by forming a single covalent bond (in this case with another F atom) Hence, we know we have a single bond in F2: F–F F2 F F This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

  21. BUT we know that the F–F molecule has 18 electrons (2 x 9) How can we arrange 18 electrons in molecular orbitals and end up with only ONE bond? • SOLUTION: • For every bonding orbital there must be an ‘anti-bonding orbital’ • An electron in a bonding orbital is cancelled out by an electron in an anti-bonding orbital This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

  22. This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

  23. Consider the MO diagram of F2 σ*(2pZ) π*(2px) π*(2py) Energy π(2px) π(2py) 2p 2p σ(2pZ) σ*(2s) 2s 2s F σ(2s) F This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

  24. Heteronuclear Diatomic molecule MO σ*(2s) Energy 2s 2s X Y σ*(2s) Homonuclear MO diagrams are symmetrical. Heteronuclear MOs are asymmetrical – the energies of equivalent atomic orbitals are DIFFERENT This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

  25. LiH molecule σ*(2s) Energy 2s 2s Li H σ*(2s) Only valence orbitals shown. The 1s (H) and 2s (Li) overlap to form the  and * molecular orbitals This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

  26. HF σ* 1s Energy Non-bonding 2p σ Non-bonding 2s HF F H The 2pz(F) can overlap with the 1s(H). The orbitals that do not overlap form NON-BONDING MOs This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

  27. Bonding 1s 2pz Anti- Bonding H F 2px 1s F H The 1s orbital on the H overlaps with the 2pz on the F to form a -bond. No overlap can occur between the 1s and the 2px or 2py, as these are pointing in the wrong direction This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

  28. HF + H–F - σ* 1s Energy Non-bonding 2p σ Non-bonding 2s HF F H The electrons are sat closer to the F atomic orbitals than the H atomic orbitals. Therefore it is predicted that the H–F bond would be POLAR This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

  29. LiF - σ* + Li–F 2s Energy Non-bonding 2p σ Non-bonding 2s LiF F Li This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

  30. Hence, the MO theory can predict POLAR bonds This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

  31. Summary This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

  32. What you should know… • Difference between homonuclear and heteronuclear bonds • Explain why some heteronuclear bonds are harder than expected to break • How the presence of hydrogen bonding in molecules affects some of their physical properties, like boiling points • How to draw the MO diagram of a heteronuclear diatomic molecule, and understand how bonding, anti-bonding and non-bonding orbitals are formed • Use the MO diagram to determine whether the bonding is likely to be polar This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

  33. Definitions… • Homonuclear bond • Heteronuclear bond • Polar bond • Hydrogen bond • Electronegativity This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

  34. Acknowledgements • JISC • HEA • Centre for Educational Research and Development • School of natural and applied sciences • School of Journalism • SirenFM • http://tango.freedesktop.org This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

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