Why does a stronger acid have a weaker conjugate base?

Let's think of it from the perspective of the Lewis and Bronsted-Lowry definitions.

Recall that:

• A Bronsted-Lowry acid donates a proton to a Bronsted-Lowry base, which accepts it.
• A Lewis acid accepts an electron pair from a Lewis base, which donated it.

LEWIS PERSPECTIVE

A strong Lewis acid would strongly want to accept an electron pair. This has to result in a bond breaking, because as an atom tries to accept valence electrons, it tries to fill its valence orbitals, and has to free up some space to accept those electrons.

This bond is usually with an #"H"^(+)#.

Therefore, a strong Lewis acid forms its conjugate Lewis base easily. Because the Lewis base forms easily, it does not strongly want to donate its electron pair (that it just received) to get a proton. That makes it a weak (Lewis) base.

BRONSTED-LOWRY PERSPECTIVE

A strong Bronsted-Lowry acid would strongly want to donate a proton to something else. That requires that the bond to that #"H"^(+)# be broken to form the conjugate base, and it follows that the bond is easily broken.

A bond that is easily broken is hard to form again. So, the conjugate Bronsted-Lowry base that forms weakly wants to accept a proton. That makes it a weak (Bronsted-Lowry) base.

From either the Lewis or Bronsted-Lowry definition, we thus come to the same conclusions:

• Weak acids have strong bonds to their #"H"^(+)#.
• Strong acids have weak bonds to their #"H"^(+)#.
• Weak bases have a strong tendency to get #"H"^(+)# and/or donate an electron pair.
• Strong bases have a weak tendency to get #"H"^(+)# and/or donate an electron pair.

You can argue in a similar manner for why the conjugate acid of a strong base is a weak acid.