Covalent Bonding - Chapter 8


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Covalent Bonding - Chapter 8

  1. 1. Covalent Bonding Or How I Learned to Love Sharing (But Remember, File Sharing is Illegal)
  2. 2. As you should remember, ionic compounds are solids at room temperatures that have one ion strip the electron(s) from the other elements’ electron cloud. Some compounds do not give up their electrons so easily. There exists a kind of tug-of-war between the two atoms. It ends in a standoff usually. This sharing of electrons is called a covalent bond . Consider hydrogen, H 2 , the simplest molecule. A hydrogen atom has a single valence electron. H• hydrogen atom + • H hydrogen atom H H hydrogen molecule • • Shared pair of electrons
  3. 3. A single covalent bond is formed when a pair of electrons is shared between two, usually nonmetal , atoms.
  4. 4. It is helpful to show the pair of electrons in a covalent bond as a dash, as in H-H for hydrogen. This is known as a structural formula. Structural formulas are the chemical formulas that show the arrangements of atoms in molecules. The dashes always indicate a pair of shared electrons, never ionic bonds. Group 4A-7A nonmetal elements are prone to form covalent bonds. + or • • F • • •• atom • • • • • F • • atom • F F • • • • • • • • • • • • • • molecule • • structural formula • • • • • • • • • • F - F Unshared pair shared pair
  5. 5. The pairs of valence electrons that are not shared between atoms are called unshared pairs. (Also called lone pair.)
  6. 6. Double and Triple Bonds Atoms sometimes share more than one pair of electrons to attain noble-gas configurations. Double covalent bonds involve two shared pairs of electrons. Two shared pairs Double bond
  7. 7. Triple covalent bonds involve three shared pairs. Electron dot formulas (Lewis structures) can also be used for more complex compounds than nitrogen or hydrogen, such as water, ammonia, methane, and carbon dioxide. Three shared pair Triple bond
  8. 8. H• H• + or Ammonia is NH 3 . H• H• H• + or Molecule (lewis structure) Molecule (structural formula) Notice the number of electrons around each atom: N = 8; stable H = 2; stable
  9. 9. Methane is CH 4 . H• H• H• H• + or Carbon dioxide is CO 2 . + or = 8 e - ’s 8 e - ’s 4 e - ’s
  10. 10. With the exception of hydrogen, all atoms in a covalent molecule must have 8 electrons around it.
  11. 11. Coordinate Covalent Bonds Most of the atoms we have seen have been fairly simple. There are some covalent bonds that don’t follow the more simple bonding rules. Carbon monoxide (CO) is an example. The carbon is four short and oxygen is two short of achieving a noble-gas configuration. A coordinate covalent bond is used. + Oxygen = 8 around it carbon = 6 around it (unstable) C O •• •• • • •• ••
  12. 12. <— = coordinate covalent bond A coordinate covalent bond is formed when one atom contributes both bonding electrons in a covalent bond. A coordinate bond is shown as an arrow that points from the atom donating the electrons. C O •• •• • • •• ••
  13. 13. Resonance Consider the two electron dot structures for ozone. Notice that the structure on the left can be converted to the one on the right by shifting electrons pairs without changing the positions of the oxygen atoms. As drawn, these electron dot structures suggest that the bonding in ozone consists of one single coordinate covalent bond and one double covalent bond.
  14. 14. Because earlier chemists imagined that the electron pairs rapidly flop back and forth, or resonate, between the various electron dot structures, they used double-headed arrows to indicate that two or more structures are in resonance . Thus the two electron dot formulas for ozone are examples of what are still referred to as resonance structures. Resonance structures are structures that occur when it is possible to write two or more valid electron dot formulas that have the same number of electron pairs for a molecule or ion.
  15. 15. Properties of Molecular Compounds Most molecular compounds are either liquids or gases, although some can be solids. Molecular compounds tend to have relatively lower melting and boiling points than ionic compounds. Molecular compounds usually don’t dissolve in water, while ionic compounds do. p.214
  16. 16. VSEPR Theory Electron dot structures fail to reflect the three-dimensional shapes of molecules. The electron dot structure and structure formula of methane (CH 4 ), for example, show the molecule as if it were flat and merely two-dimensional. Methane (electron dot structure) Methane (structural formula)
  17. 17. In reality, methane molecules are three-dimensional. The hydrogens in a methane molecule are at the four corners of a geometric solid called a regular tetrahedron. In this arrangement, all of the H-C-H angles are 109.5º, the tetrahedral angle . 109.5º
  18. 18. The valence-shell electron-pair repulsion theory, or VSEPR theory, explains the three-dimensional shape of methane. VSEPR theory states that because electron pairs repel, molecular shape adjusts so the valence-electron pairs are as far apart as possible. The methane molecule has four bonding electron pairs and no unshared pairs. The bonding pairs are farthest apart when the angle between the central carbon and its attached hydrogens is 109.5º. Any other arrangement tends to bring two bonding pairs of electrons closer together.
  19. 19. Unshared pairs of electrons are also important when you are trying to predict the shapes of molecules. The nitrogen in ammonia (NH 3 ) is surrounded by four pairs of valence electrons, so you might predict the tetrahedral angle of 109.5º for the H-N-H bond angle. However, one of the valence-electron pairs is an unshared pair. Thus they are held closer to the nitrogen than are the bonding pairs. The unshared pair strongly repels the bonding pairs, pushing them closer together than might be expected. The experimentally measured H-N-H bond angle is only 107º. This is known as the pyramidal angle .
  20. 20. 107º
  21. 21. In a water molecule, oxygen forms single covalent bonds with two hydrogen atoms. The two bonding pairs and the two unshared pairs of electrons form a tetrahedral arrangement around the central oxygen. Thus the water molecule is planar (flat) but bent. With two unshared pairs repelling the bonding pairs, the H-O-H bond angle is compressed in comparison with the H-C-H bond angle in methane. The experimentally measured bond angle in water is about 105º. This is known as the bent angle .
  22. 22. 105º
  23. 23. In contrast, the carbon in a carbon dioxide molecule has no unshared electron pairs. The double bonds joining the oxygens to the carbon are farthest apart when the O=C=O bond angle is 180º. This is known as the linear angle .
  24. 24. An exception to the octet rule is boron, B. When it forms compounds, it only needs 6 electrons (3 electron pairs) to be stable. It’s VSEPR shape is trigonal planar .
  25. 26. Polar Bonds and Molecules Covalent bonds involve electron sharing between atoms. However, covalent bonds differ in terms of how the bonded atoms share the electrons. The character of the bonds in a given molecule depends on the kind and number of atoms joined together. These features, in turn, determine the molecular properties. The bonding pairs of electrons in covalent bonds are pulled, as in a tug-of-war, between the nuclei of the atoms sharing the electrons.
  26. 27. • • +
  27. 28. When the atoms in the bond pull equally (as occurs when like atoms are bonded), the bonding electrons are shared equally, and the bond is a nonpolar covalent bond . Molecules of hydrogen (H 2 ), oxygen (O 2 ), and nitrogen (N 2 ) have nonpolar covalent bonds.
  28. 29. When a covalent bond joins two atoms of different elements and the bonding electrons are shared unequally, the bond is a polar covalent bond , or simply a polar bond . The more electronegative atoms will have the stronger electron attraction and will acquire a slightly negative charge. The less electronegative atoms will acquire a slightly positive charge.
  29. 30. Consider the hydrogen chloride molecule (HCl). Hydrogen has an electronegativity of 2.1 and chlorine has an electronegativity of 3.0. These values are significantly different, so the covalent bond is polar. H Cl
  30. 31. The chlorine atoms acquires a slightly negative charge. The hydrogen atom acquires a slightly positive charge. The lowercase Greek letter delta (  ) shows that atoms involved in the covalent bond acquire only partial charges, much less than 1+ or 1-. The minus sign in this notation shows that chlorine has acquired a slightly negative charge. The plus sign shows that hydrogen has acquired a slightly positive charge. The polarity of a bond may also be represented with an arrow pointing to the more electronegative atom.
  31. 33. The O-H bonds in the water molecule are also polar. The highly electronegative oxygen partially pulls the bonding electrons away from hydrogen. The oxygen acquires a slightly negative charge. What charge does hydrogen acquire?
  32. 34. p. 462
  33. 35. Polar Molecules The presence of a polar bond in a molecule often makes the entire molecule polar. In a polar molecule , one end of the molecule is slight negative and the other end is slightly positive. In the hydrogen chloride molecule, for example, the partial charges on the hydrogen and chlorine atoms are electrically charged regions or poles. A molecule that has two poles is called a dipolar molecule, or dipole .
  34. 36. Polar molecules orient randomly. Polar molecules line up.
  35. 37. The effect of polar bonds on the polarity of an entire molecule depends on the shape of the molecule and the orientation of the polar bonds. A carbon dioxide molecule, for example, has two polar bonds and is linear. Note that the carbon and oxygens lie along the same axis. Therefore, the bond polarities cancel because they are in opposite directions. Carbon dioxide is thus a nonpolar molecule, despite the presence of two polar bonds.
  36. 38. The water molecule also has two polar bonds. However, the water molecule is bent rather than linear. Therefore, the bond polarities do not cancel.
  37. 39. Polar vs. Nonpolar