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Bonding

Bonding. Chapter 12. 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. Energy. 0. Internuclear Distance. Energy. 0.

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Bonding

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  1. Bonding Chapter 12

  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. Energy 0 Internuclear Distance

  4. Energy 0 Internuclear Distance

  5. Energy 0 Internuclear Distance

  6. Energy 0 Internuclear Distance

  7. Energy 0 Bond Length Internuclear Distance

  8. Energy Bond Energy 0 Internuclear Distance

  9. 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.

  10. What about covalent compounds? • The electrons in each atom are attracted to the nucleus of the other. • The electrons and nuclei repel each other, • They reach a distance with the lowest possible energy. • The distance between is the bond length.

  11. Covalent Bonding • Electrons are shared by atoms. • These are two extremes. • In between are polar covalent bonds. • The electrons are not shared evenly. • One end is slightly positive, the other negative. • Indicated using small delta d.

  12. The Relationship Between Electronegativity and Bond Type

  13. 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.

  14. d+ d- H - F

  15. d+ H - F d+ H - F d- d- d- H - F d+ d+ d- H - F H - F d+ d- d- H - F d+ d+ d+ H - F H - F d- d-

  16. d+ H - F d+ H - F d- d- d- H - F d+ d+ d- H - F H - F d+ d- d- H - F d+ d+ d+ H - F H - F d- d- - +

  17. d+ d- d+ d- H - F H - F d+ d- H - F d+ d- d+ d- H - F H - F d+ d- d+ d- d+ d- H - F H - F H - F - +

  18. An Electrostatic Potential Map of HF

  19. The Charge Distribution in the Water Molecule

  20. The Carbon Dioxide Molecule

  21. Size of Isoelectronic ions • Iso - same • Iso electronic ions have the same # of electrons • Al+3 Mg+2 Na+1 Ne F-1 O-2 and N-3 • All have 10 electrons. • All have the configuration 1s22s22p6

  22. Size of Isoelectronic ions • Positive ions have more protons so they are smaller. N-3 O-2 F-1 Ne Na+1 Al+3 Mg+2

  23. Forming Ionic Compounds • Lattice energy - the energy associated with making a solid ionic compound from its gaseous ions. • M+(g) + X-(g) ® MX(s) metal nonmetal ionic compound • This is the energy that “pays” for making ionic compounds.

  24. The Structure of Lithium Fluoride

  25. Localized Electron Model • Simple model, easily applied. • A molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms. • Three Parts • Valence electrons using Lewis structures • Prediction of geometry using VSEPR • Description of the types of orbitals

  26. Exceptions to the octet rule • BH3 B wants 6 • Be and B often do not achieve octet • Have less than an octet, for electron deficient molecules. • SF6 • Third row and larger elements can exceed the octet • Use 3d orbitals? • I3-

  27. Exceptions to the octet • When we must exceed the octet, extra electrons go on central atom. • ClF3 • XeO3 • ICl4-

  28. Resonance • Sometimes there is more than one valid structure for an molecule or ion. • NO3- • Use double arrows to indicate it is the “average” of the structures. • It doesn’t switch between them; the electrons are delocalized • NO2-

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