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?

1 Answer
Dec 9, 2015

See explanation.


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.