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Possible Errors Made Regarding Polarity and Intermolecular Bonds

  1. Sep 30, 2011 #1
    Hello,

    I wouldn't necessarily regard this as homework. While these are errors I believe I have found within an online course in which I am currently enrolled, my asking is simply carried out with interest in acquiring an accurate knowledge of entry-level chemistry.

    While studying polar molecules and various intermolecular forces, such as the London force, dipole-dipole interactions and hydrogen bonds, I came across a sample problem that appeared to contain some inaccuracies.

    In one section of the sample, my resource claims that Arsine (AsH3) is a polar molecule. This doesn't make sense to me. While the structure of Arsine isn't symmetrical, both hydrogen and arsenic have an electronegativity of 2.2. Would this not imply that the molecule is nonpolar, as no polarity due to a difference in the electronegativity of the bonded atoms is established?

    Also, within the same sample problem, my resource claims that hydrogen fluoride (HF) only contains hydrogen bonds and that hydrogen iodide (HI) only contains dipole-dipole interactions. While the aforementioned intermolecular interactions do occur within these two molecules respectively, can HF not also contain dipole-dipole interactions and the London force? Can HI not contain the London force in addition to its dipole-dipole interactions? Recognizing that only one type of intermolecular interaction is listed for each, I ask: is it common practice to only list the interaction that affects a molecule the most?

    I thank anyone who has taken the time to read and consider my concerns, your contributions are greatly appreciated.

    Thank you,

    Eric.
     
  2. jcsd
  3. Sep 30, 2011 #2
    Arsine has an electron lone pair, correct? Do you think that could possibly give rise to a dipole moment? Remember, a dipole moment exists when you have a non-uniform charge distribution.

    Generally, the focus is on the predominant interaction that governs intermolecular behavior. Hydrogen bonding has been termed by some to be a very strong dipole-dipole interaction (the details can get a bit involved, and the topic of countless research over the years). London dispersion forces arise from an instantaneous dipole inducing a transient dipole moment in a neighboring molecule. While dispersion forces are always present, they might be very small if one is dealing a molecule with a very strong permanent dipole moment. Generally, the strength of London dispersion forces will increase with larger molecules, as they become more polarizable.

    Certainly, though, chemical (and biochemical) interactions are the sum of many interactions. The question is whether or not all of those effects are actually present in the system at hand, and if so, how strong are they.
     
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