Question

If HSO3- and HC2O4- are both put into water what will the equilibrium equation be?

We are told to ignore reactions between the ions and water.

The only thing I am confused by is that both the substances are amphoteric and both predominately act as acids in water so how do I write the products?

Answer 1

`HSO_3^(-) +HC_2O_4^(-) rightleftharpoonsSO_3^(2-)+C_2O_4H_2(aq)`

Explanation:

...else...

`HSO_3^(-) +HC_2O_4^(-) rightleftharpoonsH_2SO_3(aq)+C_2O_4^(2-)`

Where will the equilibrium lie? I dunno. And unless you quote data for the dissociation constants of oxalic and sulphurous acid you don't know either.

Answer 2

Here's what I get.

Explanation:

Both substances are acids, and both are amphoteric.

When both of them are together, the stronger acid will protonate the weaker acid, which will act as a Brønsted base.

Which is the stronger acid?

We have the following equilibria:

`"HSO"_3^"-" + "H"_2"O" ⇌ "H"_3"O"^"+" + "SO"_3^"2-"; color(white)(ml)K_text(a₂) = 6.2 ×10^"-8"`

`"HC"_2"O"_4^"-" + "H"_2"O" ⇌ "H"_3"O"^"+" + "C"_2"O"_4^"2-"; K_text(a₂) = 6.5 ×10^"-5"`

The hydrogen oxalate ion is the stronger acid. It will protonate the hydrogen sulfite ion.

Write the equation for the equilibrium

`overbrace(underbrace("HC"_2"O"_4^"-")_color(red)("stronger acid"))^(color(brown)("Brønsted acid")) + overbrace(underbrace("HSO"_3^"-")_color(red)("weaker acid"))^(color(brown)("Brønsted base")) ⇌ "C"_2"O"_4^"2-" + "H"_2"SO"_3^"-"`