Question #e09dc

By the Bronsted-Lowry definition:

• A Bronsted acid donates protons.
• A Bronsted base accepts protons.

#"HPO"_4^(2-) + "H"_2"O" rightleftharpoons "H"_2"PO"_4^(-) + "OH"^(-)#

Note that #"HPO"_4^(2-)# gained a proton to become #"H"_2"PO"_4^(-)#. Thus it is a Bronsted base and its conjugate acid is #"H"_2"PO"_4^(-)#.

Note that #"H"_2"O"# lost a proton to become #"OH"^(-)#. Thus it is a Bronsted acid and its conjugate base is #"OH"^(-)#.

#color(blue)(stackrel("Bronsted Base")overbrace("HPO"_4^(2-)) + stackrel("Bronsted Acid")overbrace("H"_2"O") rightleftharpoons stackrel("Conjugate Acid")overbrace("H"_2"PO"_4^(-)) + stackrel("Conjugate Base")overbrace("OH"^(-)))#

• The Bronsted base and the conjugate acid are one of the two acid-base pairs.
• The Bronsted acid and the conjugate base are the other acid-base pair.

If you compare this back to the Lewis definition:

• A Lewis acid accepts electron pairs.
• A Lewis base donates electron pairs.

then an interesting connection between the two is that:

• In acid-base reactions, electrons are often accepted by a Lewis acid in exchange for a donated proton that is accepted by a Bronsted base.
• In acid-base reactions, electrons are often donated by a Lewis base in exchange for an accepted proton donated from a Bronsted acid.

Thus:

• A Lewis base that becomes protonated is also a Bronsted base.
• A Lewis acid that becomes deprotonated is also a Bronsted acid.