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Trouble with picturing the 3-d structures

  1. Jun 1, 2006 #1
    :surprised I know I am overthinking, as always, but for some reason I am just having so much trouble with picturing the 3-d structures (dash/wedge) when solving the problems. My test is tomorrow and I am stuck, I just can't draw them correctly...I get the hybridization, and the angle that will form.
    What I don't get is how do you know where to draw a dash or a wedge, and when will the dash be to the upper left corner (extended) versus the lower left? Sometimes they are drawn near eachother (dash and wedge going towards the upper left) with the one in the plane going towards the left. How do you determine this? Does it have to do with the bond angles? I tried figuring it out that way, but I still can't get the right figure without looking at the answer for a hint. For example: (CH3)3 N ... I have no idea how to orient my lewis dot to come up with the right three-d...how do you know which of the hydrogens will be off the plane and et.? Is there a simple way of doing this, and what is wrong with the way I am thinking?
    Also, for a positive charged [H2COH] (+) molecule, why is the oxygen atom (as it says in my book) sp2? I know the carbon is sp2, but wouldn't trhe oxygen be sp3 because it's bonded to two things, and it also has 2 pairs of lone electrons?
    Any help is greatly appreciated.

    Thank you for your help in advance!
     
  2. jcsd
  3. Jul 11, 2006 #2
    To extend the ideas of valence-bond theory to polyatomic molecules, we must envision mixing s, p, and sometimes d orbitals to form hybrid orbitals. The process of hybridization leads to hybrid atomic orbitals that have a large lobe directed to overlap with orbitals on another atom to make a bond. Hybrid orbitals can also accommodate nonbonding pairs. A particular mode of hybridization can be associated with each of the five common electron-domain geometries ( trigonal trigonal and ).

    Covalent bonds in which the electron density lies along the line connecting the atoms (the internuclear axis) are called sigma bonds. Bonds can also be formed from the sideways overlap of p orbitals. Such a bond is called a pi bond. A double bond, such as that in consists of one bond and one bond; a triple bond, such as that in consists of one and two bonds. The formation of a bond requires that molecules adopt a specific orientation; the two groups in for example, must lie in the same plane. As a result, the presence of bonds introduces rigidity into molecules. In molecules that have multiple bonds and more than one resonance structure, such as the bonds are delocalized; that is, the bonds are spread among several atoms.
     
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