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Comparing acidity among 3 compounds - pic included

  1. Oct 31, 2008 #1
    1. The problem statement, all variables and given/known data
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    My book says that if I attach an electron-donating group to an acid, it would decrease the acidity of that acid because the electrons would cause the already negative ion to be more unstable.

    Well, -OH is an electron-donating group. I don't understand why when I attach an -OH group to acetic acid to make it into 2-hydroxyacetic acid, its acidity actually increases. It is suppose to decrease, according to the book. Does anybody know how to resolve this conflict?

    Also, if I compare phosphoric acid with carboxylic acid, the phosphoric acid is a lot stronger than carboxylic acid. I don't understand why. If I deprotonate 1 acid, the negative charge would form resonance with one other oxygen on both acids (phosphoric acid and carboxylic acid). I know phosphoric acid has 4 oxygens attached, but two of them would probably not be involved in resonance stabilization because they are too busy bonding to Phosphorus and Hydrogen. So, if both phosphoric acid and carboxylic acid can only form 2 resonance structures when 1 H is deprotonated, then why is phosphoric acid so much stronger than carboxylic acid?

    Thanks in advance.
     
    Last edited: Oct 31, 2008
  2. jcsd
  3. Nov 1, 2008 #2
    i would use the inductive effect to explain the acidity of acetic acid and 2 hydroxy acetic acid. the oxygen on the hydroxyl group of 2 hydroxy acetic acid is of course electronegative. it tends to pull electrons towards itself. the carbon of -COOH will have a lesser electron density, and the O-H bond of the -COOH is weakened that is the hydrogen is held less strongly and is easily removed.

    i think it is like when you attach a chlorine atom instead of the -OH.
     
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