Bonding and Molecular Structure: Fundamental Concepts Valence e- and Bonding Covalent Ionic - PowerPoint PPT Presentation

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Bonding and Molecular Structure: Fundamental Concepts Valence e- and Bonding Covalent Ionic

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  1. AP Notes Chapter 8 Bonding and Molecular Structure: Fundamental Concepts Valence e- and Bonding Covalent Ionic Resonance & Exceptions to Octet Rule Bond Energy & Length Structure, Shape & Polarity of Compounds

  2. What is a Bond? • A force that holds atoms together. • Why? • We will look at it in terms of energy. • Bond energy the energy required to break a bond. • Why are compounds formed? • Because it gives the system the lowest energy.

  3. Covalent compounds? • The electrons in each atom are attracted to the nucleus of the other. • The electrons repel each other, • The nuclei repel each other. • The reach a distance with the lowest possible energy. • The distance between is the bond length.

  4. Thus Hydrogen is Diatomic! Bond Formation

  5. e- Covalent Character

  6. He2 . . E He + He . . Inter-nuclear Distance Why Isn’t Helium Diatomic?

  7. F + F F2 2p____ ____ ___ ___ ____ ____ 2p2s ____ ____ 2s F F

  8. Ionic Bonding • An atom with a low ionization energy reacts with an atom with high electron affinity. • The electron moves. • Opposite charges hold the atoms together.

  9. Li + Cl1s22s1 [Ne] 3s23p52s ___ 3p _____ _____ ___1s _____ 3s _____[Ne]

  10. Li + Cl 2s ___ 3P _____ _____ _____1s _____ 3s _____[Ne]

  11. LiCl2s ___3P _____ _____ _____1s _____ 3s _____ [Ne]

  12. Electronegativity Describes the relative ability of an atom within a molecule to attract a shared pair of electrons to itself.

  13. Electronegativity Pauling electronegativity values, which are unit-less, are the norm.

  14. ElectronegativityRange from 0.7 to 4.0 Figure 9.9 – Kotz & Treichel

  15. Bond: A - B DEN = | ENA - ENB |

  16. Bond Character “Ionic Bond” - Principally Ionic Character “Covalent Bond” - Principally Covalent Character

  17. covalent ionic EN ~0 1.7 ~4 Determining Principal Character of Bond

  18. F - F EN = 0 Non-polar

  19. N - O EN = |3.0 - 3.5| = 0.5 O N Slightly polar

  20. Ca - O  EN = |1.0 - 3.5| = 2.5 Ca O Ionic Bond with somecovalent character

  21. Electronegativity • The ability of an electron to attract shared electrons to itself. • Pauling method • Imaginary molecule HX • Expected H-X energy = H-H energy + X-X energy 2 • D = (H-X) actual - (H-X)expected

  22. Electronegativity • D is known for almost every element • Gives us relative electronegativities of all elements. • Tends to increase left to right. • decreases as you go down a group. • Noble gases aren’t discussed. • Difference in electronegativity between atoms tells us how polar.

  23. Polar Covalent Ionic Electronegativity difference Bond Type Zero Covalent Covalent Character decreases Ionic Character increases Intermediate Large

  24. Dipole Moments • A molecule with a center of negative charge and a center of positive charge is dipolar (two poles), • or has a dipole moment. • Center of charge doesn’t have to be on an atom. • Will line up in the presence of an electric field.

  25. d+ d- H - F How It is drawn

  26. Which Molecules Have Them? • Any two atom molecule with a polar bond. • With three or more atoms there are two considerations. • There must be a polar bond. • Geometry can’t cancel it out.

  27. Ionic Radii -- Cations

  28. Ionic Radii -- Anions

  29. Molecular Polarity Vector Sum of Bond Polarities

  30. MgBr2 Mg - Br EN = |1.2 - 2.8| = 1.6 Mg Br Br Covalent BOND w/much ionic character, BUT NON-POLAR molecule

  31. Lewis Structures

  32. The most important requirement for the formation of a stable compound is that the atoms achieve noble gas e- configuration

  33. Valence Shell ElectronPair Repulsion Model(VSEPR) The structure around a given atom is determined principally by minimizing electron-pair repulsions

  34. Electron pairs Bond Angles Underlying Shape 2 180° Linear 3 120° Trigonal Planar 4 109.5° Tetrahedral 90° & 120° Trigonal Bipyramidal 5 6 90° Octagonal VSEPR

  35. LEWIS STRUCTURES • : draw skeleton of species • : count e- in species • : subtract 2 e- for each bond in skeleton • : distribute remaining e-

  36. Distinguish Between ELECTRONIC Geometry & MOLECULAR Geometry

  37. CH4 Bond angle = 109.50 Electronic geometry: tetrahedral Molecular geometry: tetrahedral

  38. H3O+ Bond angle ~ 1070 Electronic geometry: tetrahedral Molecular geometry: trigonal pyramidal

  39. H2O Bond angle ~ 104.50 Electronic geometry: tetrahedral Molecular geometry: bent

  40. NH2- Bond angle ~ 104.50 Electronic geometry: tetrahedral Molecular geometry: bent

  41. “Octet Rule” holds for connecting atoms, but may not for the central atom.