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Bonding and Molecular Structure: Valence e- and Bonding Covalent Ionic Bond Energy & Length

Learn about the forces that hold atoms together, bond energy, covalent and ionic bonds, bond length, structure, shape, and polarity of compounds. Explore differences in electronegativity, bond character, and dipole moments. Understand Lewis structures and VSEPR theory to predict molecular geometries.

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Bonding and Molecular Structure: Valence e- and Bonding Covalent Ionic Bond Energy & Length

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  1. Chapter 6 Chemical Bonding Bonding and Molecular Structure: Valence e- and Bonding Covalent Ionic 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 The difference between ionic and covalent bonds. 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

  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 • 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.

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

  23. 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.

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

  25. 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.

  26. Ionic Radii -- Cations

  27. Ionic Radii -- Anions

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

  29. Lewis Structures

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

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

  32. 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

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

  34. Distinguish Between ELECTRONIC Geometry & MOLECULAR Geometry

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

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

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

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

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

  40. BaI2 Bond angle =1800 Electronic geometry: linear Molecular geometry: linear

  41. BF3 Bond angle =1200 Electronic geometry: trigonal planar Molecular geometry: trigonal planar

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