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Chemistry 16: 9-10

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Acids and Bases: Lewis Acids


Lewis Acids and Bases

We have two definitions of acids and bases so far:

Types of Acids

We can now make some general rules about shapes of molecules, chemical bonds, and acid reaction properties.

Halide acids HXIf the halide bond is strong, the acid will not tend to dissociate, and the acid will be weak. The relative strength depends on the bond strength plus the electron attachment enthalpy, with bond strength being the determining factor. The greater this sum, the weaker the acid.
Oxoacids: non-metal atoms attached to oxygen, some of which is attached to hydrogen (e.g. HNO2)Again, acid strength depends on bond strength plus electron affinity, but in this case, electron affinity is the determining factor. Adding oxygen to the complex distributes the negative charge so that the H is more easily dissociated, making HNO3 a stronger acid than HNO2.
Carboxylic Acids: acids containing -CO2H groupsThe addition of electronegative substitutes for one of the H atoms increases the acidity of a compound: CH3CO2H (acetic acid) is less acidic than ClCH2CO2H (chloroacetic acid).
Hydrated Metal Acids (M+ • nH2O)Because the metal ion M+ is positively charged, the attracted water molecule are more extremely polarized and the H+ ion is easily detached, making the complex as a whole a proton donor.
Anion Bases (negatively charged atoms or molecules) The negative charge on the anion attracts H atoms in water molecules, causing them to dissociate and form bonds with the anion, so the anion acts as a base, turning water into an acid proton donor.

Practice with the Concepts

  • Can water be a Brønsted-Lowry acid?
  • Can water be a Brønsted-Lowry base?
  • Can water be a Lewis acid?
  • Can water be a Lewis base?

Discussion Questions

Optional Readings

You may want to revisit the ChemGuide page on Acids and Bases for more examples.