The hydrogen oxalate ion is amphoteric. Write two balanced chemical equations to illustrate this property. What is the conjugate acid of this ion and what is its conjugate base?

As you know, an amphoteric compound can act both as a Bronsted - Lowry acid and a Bronsted - Lowry base.

More specifically, the hydrogen oxalate ion, #"HC"_2"O"_4^(-)#, will act as an acid by donating a proton, #"H"^(+)#, and as a base by accepting a proton.

Now, what would you expect to see when the hydrogen oxalate ion acts as an acid in aqueous solution?

Well, since it donates a proton, its net charge will go from #1-# to #2-#, since a proton carries a #1+# charge.

Moreover, the acidity of the solution should increase, i.e. its pH should decrease. This tells you that the concentration of hydronium ions, #"H"_3"O"^(+)#, should Increase.

#"CH"_2"O"_text(4(aq])^(-) + "H"_2"O"_text((l]) rightleftharpoons "C"_2"O"_text(4(aq])^(2-) + "H"_3"O"_text((aq])^(+)#

The conjugate base of a Bronsted - Lowry acid is the chemical species that remains after the acid donates a proton.

In this case, the oxalate anion, #"C"_2"O"_4^(2-)#, will be the hydrogen oxalate ion's conjugate base.

How about when the hydrogen oxalate acts as a base in aqueous solution?

Well, this time it's accepting a proton, so its charge will go from #1-# to #0#.

Moreover, the basicity of the solution will Increase, i.e. its pH will increase. This tells you that teh concentration of hydroxide ions, #"OH"^(-)#, will increase.

#"CH"_2"O"_text(4(aq])^(-) + "H"_2"O"_text((l]) rightleftharpoons "H"_2"C"_2"O"_text(4(aq]) + "OH"_text((aq])^(-)#

The conjugate acid of a Bronsted - Lowry base is the chemical species that remains after the base accepts a proton.

In this case, oxalic acid, #"H"_2"C"_2"O"_4#, is the hydrogen oxalate's conjugate acid.