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6.S: Acid-base Reactions (Summary)

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    242881
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    Before you move on to the next chapter, you should:

    • Know the Bronsted-Lowry definition of acidity and basicity: a Bronsted acid is a proton donor, a Bronsted base is a proton acceptor.
    • Know the Lewis definition of acidity and basicity: a Lewis acid is an electron acceptor, a Lewis base is an electron donor.
    • Understand that the Lewis definition is broader: all Bronsted acids are also Lewis acids, but not all Lewis acids are also Bronsted acids.
    • Be able to draw a curved arrow mechanism for both Bronsted and Lewis acid-base reactions.
    • Know the expressions for \(K_a\) and \(pK_a\).
    • Commit to memory the approximate pKa values for the following functional groups:
      • \(H_3O^+\), protonated alcohol, protonated carbonyl (~ 0)
      • carboxylic acids (~ 5)
      • imines (~ 7)
      • protonated amines, phenols, thiols (~ 10)
      • water, alcohols (~ 15)
      • \(\alpha \)-carbon acids (~ 20)
    • Be able to use \(pK_a\) values to compare acidity: a lower \(pK_a\) corresponds to a stronger acid.
    • Know that:
      • For a given pair of acids, the stronger acid will have the weaker conjugate base.
      • For a given pair of basic compounds, the stronger base will have the weaker conjugate acid.
    • Be able to identify the most acidic/basic groups on a polyfunctional molecule.
    • Be able to calculate the equilibrium constant of an acid base equation from the \(pK_a\) values of the acids on either side of the equation.
    • Be able to use the Henderson-Hasselbalch equation to determine the protonation state/charge of an organic compound in an aqueous buffer of a given \(pH\).
    • Understand the idea that the best way to compare the strength of two acids is to compare the stability of their conjugate bases: the more stable (weaker) the conjugate base, the stronger the acid.
    • Be able to compare the acidity or basicity of compounds based on periodic trends:
      • acidity increases left to right on the table, so alcohols are more acidic than amines
      • acidity increases top to bottom on the table, so a thiol is more acidic than an alcohol.
    • Be able to compare the acidity or basicity of compounds based on protonation state: \(H_3O^+\) is more acidic than \(H_2O\), \(NH_4^+\) is more acidic than \(NH_3\).
    • Understand how the inductive effect exerted by electronegative groups influences acidity.
    • Understand how resonance delocalization of electron density influences acidity.
    • Be able to explain/predict how orbital hybridization affects the relative acidity of terminal alkynes, alkenes, and alkanes.
    • Be able to explain why phenols are more acidic than alcohols, and how electron-withdrawing or donating groups influence the acidity of phenols.
    • Be able to identify the relative basicity of a nitrogen-containg group in a compound, based on whether it is an amine, amide, imine, aniline, or 'pyrrole-like'.
    • Be able to identify -carbon(s) on a carbonyl compound, and explain why protons are weakly acidic. You should be able to draw the enolate conjugate base of a carbonyl compound.
    • Be able to identify tautomeric relationships, specifically keto-enol and imine-enamine tautomers.
    • Understand what a polyprotic acid is, what is meant by multiple pKa values, and why these values get progressively higher.

    Contributors and Attributions


    6.S: Acid-base Reactions (Summary) is shared under a CC BY-NC-SA license and was authored, remixed, and/or curated by Tim Soderberg.

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