# Which species dissociate completely in ionic equations?

I am tasked with writing a molecular equation, an overall ionic equation, and a net ionic equation for the reaction between an aqueous solution of phosphoric acid, $\ce{H3PO4{(aq)}}$ and an aqueous solution of sodium hydroxide. The products are sodium phosphate and water.

I have already figured out the balanced molecular equation of the following:

$$\ce{H3PO4 {(aq)} + 3 NaOH {(aq)} -> Na3PO4 {(aq)} + 3 H2O {(l)}}$$

However, I am uncertain how to write the proper ionic equations. I know that the ions are separated, but how do you format this properly when writing with multiple molecules. If this is based on the disassociation of the chemicals, how do you determine which ions disassociate?

I have arrived at:

$$\ce{3H+ {(aq)} + PO4^3- {(aq)} + 3 Na+ {(aq)} + 3 OH- -> 3 Na+ {(aq)} + 3 H2O {(l)} + PO4^3- {(aq)}}$$

However, I still do not know why this is the best answer. Why does the $\ce{NaOH}$ show separated in the equation while the $\ce{PO4}$ does not? How would you know which ones separate by simply looking at the problem?

• Okay, I think I might somewhat understand it now. The difference is whether something is a strong acid or a weak acid. A strong acid will disassociate completely, therefore, it will be split apart completely within the equation. Other items, including weak acids, will not be split. Please let me know if my reasoning is off on this. – Jonathan Hickman Sep 29 '12 at 15:52

The second equation you’ve given appears to be the right overall ionic equation. Phosphate is a weak acid, $\ce{NaOH}$ is a strong base. We expect NaOH to dissociate at any pH, where the speciation of phosphate ($\ce{H3PO4}$, $\ce{H2PO4-}$, $\ce{HPO4^2-}$, and $\ce{PO4^3-}$) depends on pH.

A “net ionic equation” doesn’t include spectator ions. Spectator ions exist on both sides of the arrow and can be “canceled like” (think as in algebra). For any acid base reaction, it’s really easy to remember!

$$\ce{H+ {(aq)} + OH- {(aq)} <=> H2O {(l)}}$$

Lastly don’t forget the state: $\ce{OH {(aq)}-}$. I was taught to right them as subscripts, but I’ve seen then as regular text too: $\ce{OH- (aq)}$.

• If phosphate is a weak acid, it should not disassociate completely, right? So, the H3PO4 should stay together I would think. If that is the case, then should I really be breaking down the ionic equation as follows: $3HPO_4(aq)+3Na^+(aq)+3OH^-(aq) \longrightarrow 3Na^+(aq) + 3H_2O(l)+PO_4^{3-}(aq)$? – Jonathan Hickman Oct 6 '12 at 16:37

Firstly what you have here is an acid base problem. The way you can tell when you have an acid base problem is when your aqueous reactants produce water.

Secondly you know that anything with an (aq) subscript are ions in solution. AKA they are telling you which molecules are disassociated in water.