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Bonding Theory I

Bonding Theory I. Chemical bond: an interaction among electrons that results in an attraction between two or more atoms to form a compound. Bonding Theory I. What factors will effect bonding? What should be the guiding principle that results in whether atoms will bond together?.

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Bonding Theory I

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  1. Bonding Theory I • Chemical bond: an interaction among electrons that results in an attraction between two or more atoms to form a compound.

  2. Bonding Theory I • What factors will effect bonding? • What should be the guiding principle that results in whether atoms will bond together?

  3. Bonding Theory I • Electrons are involved in bonding, so…. Where are the electrons? What is the charge density?

  4. Bonding Theory I • Nature tends toward stability: atoms bond together in an effort to establish a stable electron arrangement (octet rule?).

  5. Bonding Theory I • Nature does whatever is easiest: the mechanism by which bonds form is a result of energy requirements.

  6. Bonding Theory I • Electron affinity: the ability of an atom to attract an electron to itself.

  7. Bonding Theory I • Ionization energy: the energy required to remove an electron from a gaseous atom. • What are the periodic trends of electron affinity and ionization energy?

  8. Bonding Theory I • Electronegativity: the relative ability of an atom to attract an electron when it is bonded to another atom.

  9. Bonding Theory I • Electronegativity increases across a period and decreases down a group. Why?

  10. Bonding Theory I • Polar vs Non Polar Bonds • Electronegativity describes where an electron will be as a result of bonding. Ex. F-F bond, vs. Li-F bond, vs. C-H bond

  11. Molecular Polarity • Can a molecule formed by polar bonds be non-polar? • Vector Sum of Arrows

  12. Molecular Polarity • HCl molecules line up in an electric field, Cl2 molecules do not. • Determine which are polar/non polar: BeF2, H2O, CH4, and CH3Cl

  13. Bonding Theory I • The most reactive elements are those that have the highest and lowest electronegativities.

  14. Bonding Theory I • Bond strength: a measurement of the relative attraction between bonded atoms as determined by the energy required to break the bond.

  15. Bonding Theory I • The larger the difference in electronegativity, the stronger the bond. • Which is stronger? C-H or H-F?

  16. Bond Character • Two major types of bonds: ionic and covalent. • Ionic results from large differences in negativity. • Covalent from small differences. (polar?)

  17. Bond Character • Bonds are rarely purely ionic or covalent, but rather lie on a continuum between the two: ex. only those that have no difference in electronegativity are 100% covalent.

  18. Ionic Bonds • Ionic bonds occur between atoms which have large differences in electronegativity(>1.66). Usually involves a metal attracted to a non-metal.

  19. Ionic Bonds • Ionic compounds have high melting points. They form crystal lattices and conduct electricity in their molten (melted) state.

  20. Ionic Bonds • 1.) A metal comes into close proximity to a non-metal. • 2.) One or more electrons are completely transferred to the non-metal. (why?)

  21. Ionic Bonds • 3.) The metal becomes positively charged (cation), the non-metal becomes negatively charged (anion).

  22. Ionic Bonds • An electrostatic attraction results in the ionic bond.

  23. Covalent Bonds • Occurs between non-metals which have similar electronegativities.

  24. Covalent Bonds • They have low melting points, do not conduct electricity, and are brittle. • They form molecules of varying bond lengths and angles.

  25. Covalent Bonds • Bond length: the distance between two nuclei along a bond axis. • Bond angle: the angle between two bond axes.

  26. Covalent Bonds • The bond is formed as the atoms get close together and overlap orbitals. • Electrons become located in the combined orbital.

  27. Covalent Bonds • In this common orbital, electrons become attracted to both nuclei at the same time. The electrons are considered shared between the two atoms.

  28. Polyatomic Ions • Polyatomic Ions: non-metals bond together in such a way that there is a net remaining charge. • Ex. ammonium (Dr. Cakir’s hair analogy)

  29. Bonding Theory II • BONDS INNNNN SPAAAAAAACE! • The behavior of molecules is based on the structure of its bonding.

  30. Bonding Theory II • Lewis structures show the distribution of valence electrons in an atom, molecule, or polyatomic ion. Dots = unshared electrons, Bonds = straight lines.

  31. Bonding Theory II • Lewis Dot Diagrams: • Valence electrons arranged to show how each atom fills its outer shell.

  32. Bonding Theory II • Lewis Dot Diagrams. • For all atoms that form covalent bonds, except hydrogen, 8 electrons represents a full outer level.

  33. Bonding Theory II • Lewis Dot Diagrams. • Shared pairs - 2 electrons shared in bonding. • Unshared pairs - 2 electrons not involved in bonding.

  34. Bonding Theory II • 1a.) H2Te • 1c.) NI3 • 1d.) CBr4

  35. Bonding Theory II • Rules for writing Lewis structures: • 1. Count the number of valence electrons, accounting for any charges.

  36. Bonding Theory II • 2. Draw the skeletal structure, using single bonds. • 3. Deduct two electrons for each single bond.

  37. Bonding Theory II • 4. Distribute the leftover electrons to give each atom a noble gas structure (if possible).

  38. Bonding Theory II • 5. If too few electrons: form multiple bonds. • 6. If too many electrons: expand the octet or look for octet rule violations.

  39. Bonding Theory II • Ex. OCl- , CH3OH, SO32-. • Ex. NO3-, XeF4, BF3

  40. Electron Pair Repulsion • Electron pairs spread out as far as possible to minimize repulsive forces. • Pauli - only 2 electrons of opposite spins may share the same orbital.

  41. Electron Pair Repulsion • Parallel spinning electrons may not. They repel. In fact, the electron-electron repulsion formed by different charge clouds is greater than the normal electrostatic repulsion.

  42. Electron Pair Repulsion • Thus, unshared pairs will alter molecular shape and bond angles.

  43. Electron Pair Repulsion • Unshared-unshared repulsion > unshared-shared repulsion > shared-shared repulsion > 2 electrons in the same orbital.

  44. Electron Pair Repulsion • Bond angles with no unshared pairs: • linear, trigonal planar, and tetrahedral molecules • Other shapes (like trigonal bipyramidal, octahedral)

  45. Electron Pair Repulsion • Bond angles with unshared pairs: • linear molecules • Trigonal pyramidal • Bent molecules

  46. Electron Pair Repulsion • Bonds involving atoms with unshared pairs, will result in angles that are less than expected.

  47. Electron Pair Repulsion • See table 7.3 and figure 7.8 • Note: multiple bonding has no effect on molecular geometry.

  48. Hybrid Orbitals • Atoms form hybrid orbitals before covalent bonding. • The s and p orbitals blend to form a hybrid orbital. • sp3 , sp2, sp

  49. Hybrid Orbitals • sp3: used in C-C single bonds (angle = 109.5) • sp2: used in double bonds (angle = 120) • sp: used in triple bonds (angle = 180)

  50. Sigma and Pi bonds • Sigma bond - the overlap of two orbitals that lie directly on the same bond axis. • Single bonds are composed of a sigma type bond.

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