Question

Write a net ionic equation for the reaction between H3PO4(aq) and CO32-(aq) that shows H3PO4(aq) behaving as a Bronsted-Lowry acid.

(2) Decide which would be favored at equilibrium for this reaction, reactants or products?

Answer 1

First, let us define Bronsted Lowry acid/base:

Bronsted Lowry acid: any species that will donate H+ (protons) in solution, and makes pH lower (i.e HCl)

Bronsted Lowry base: any species that will accept H+ (protons) in solution, and makes pH higher (NH3 will accept H+ to form NH4+)

Typically, acid/bases are shown in the left (reactants)

when we write the products:

Bronsted Lowery conjugate base = the base formed when the B.L. acid donates its H+ proton ( i.e. HCl -> Cl-

Bronsted Lowery conjugate acid = the acid formed when the B.L. base accept its H+ proton ( i.e. NH4+ has accept H+ proton)

Note that, typically conjugate bases/acids are shown in the right (product) side

So, from your reaction:

H3PO4(aq) = acid

CO3-2(aq)= base

then

H3PO4(aq) + CO3-2(aq) = HCO3-(aq) + H2PO4-(aq)

this is very likely to happen (forward)

since H3PO4 is stronger acid than HCO3- , then expect CO3-2 to act as a base