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Thread: Re: Acidic nature of alcohols

  1. #1 Re: Acidic nature of alcohols 
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    My original question:How does acidic nature of alcohols vary as we increase the no. of alkyl (R) groups attached to the O-H group?

    What I was able to figure out:
    As the oxygen atom is more electronegative than hydrogen, the OH bond becomes polarised. As we increase the no. of R groups, the magnitude of +I effect increases too. Consequently, electron density over oxygen increases and the electron pair of O-H bond shifts a little more towards hydrogen. Thus, the polarisation of OH bond weakens. More the no. of R groups, the weaker the polarisation gets.

    Follow-up questions:
    Did I get this right so far? If yes, then please try explaining how does the strength of polarisation determines the acidic nature of alcohols? Like if the bond is more polar, would that increase the chances of breakage of the OH bond? Or would it strenghten the bond and save it from breakage?

    Point of confusion: The latter portion of the above section is where I am getting stuck.

    Any rectification of my original explantion, any links dealing with the subject will be highly appreciated.

    Thank you


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  3. #2 Re: Acidic nature of alcohols 
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    Quote Originally Posted by spider
    how does the strength of polarisation determines the acidic nature of alcohols?
    -a more acid molecule can more easily loose it’s H.
    -a more polarised bound is a bound is which the two electron binding the two atoms are more concentrated/localized toward one of the two atoms (closer to the O further away from the Hin this case).
    -The result of the acid reaction is O- (having the two electrons between O and H for himself) en H+ (who lost his electron).
    So a more polarised bound is already closer end result of an acid reaction.

    Makes sense?


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  4. #3  
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    The +I you describe is called inductive effect, it means that an alkyl pushes it electrons away from itself thus the O receives a negative charge from the R (here = alkyl groups).
    The larger the R group result (basically) in a larger +I effect
    As a consequence the O can less easily accept the other electron of the O-H bound => making it less acid.
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  5. #4  
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    Okay HIM thank you.

    I see it now. Just one more point.

    Does all that boil down to the point that if the polarisation is increased it will enhance the chance of the cleavage of the bond (OH bond I mean)?
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  6. #5  
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    the O-H is not cleaved it is completely pulled toward the O and away from the H. So any effect which reduces the potential of the O to attract the bound to itself will reduce the acidity
    While every effect which increases the potential of the O to attract the bound will increase the acidity.

    so the polarisation results that O can more easily attract the O-H bound. But okay if wont to call this cleavage fine by me.
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  7. #6  
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    CH3COOH -- CH3COO- + H+

    Do you mean to say that the OH bond hasn't broken up in this case?
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  8. #7  
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    You can call it broken up.
    But again the bound is completely at the O
    I just told it wasn’t cleaved to let you realise that the electrons go to one side and are not divide (cleaved) over two (the H and the O).
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  9. #8  
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    Okay, thanks again. Now, I get it. Still, I would check it up with my prof.
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  10. #9  
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    Like Him said, it is about inductive electron donating effects.

    Acidity is measured by the stability of the conjugate base (the anion left behind after the proton has been donated). As has already been said, alkyl groups are electron donating which means that when the O-H bond undergoes heterolytic fission (both bonding electrons go onto one of the atoms, in this case the oxygen - compared to homolytic fission where 1 electron goes to each atom, leaving radicals) then the negative charge is heavily localised on the O. This is less stable than if you had a load of electron withdrawing groups like Cl on the molecule, which serve to spread the negative charge out a bit more - more stable anion means e.g. 2-chloroethanol is a stronger acid than ethanol.
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