# Why did they ‘split’ this H2+ ion in the ionic equation?

I am asked this chemistry question in my textbook:

Write an ionic equation for this reaction: $$\ce{Mg(s) + H2SO4(aq) -> MgSO4(aq) + H2(g)}$$

This is my thought process and answer: $$\require{cancel}$$ \begin{align} \ce{Mg(s) + H2SO4(aq) &-> MgSO4(aq) + H2(g)}\\ \ce{Mg(s) + H2+(aq) + SO4^{2-}(aq) &-> Mg^{2+}(aq) + SO4 ^{2-}(aq) + H2(g)}\\ \ce{Mg(s) + H2+(aq) + \cancel{\ce{SO4^{2-}(aq)}} &-> Mg^{2+}(aq) + \cancel{\ce{SO4^2-(aq)}} + H2(g)}\\\hline \ce{Mg(s) + H2+(aq) &-> Mg^{2+}(aq) + H2(g)} \end{align}

$$\ce{Mg(s) + 2H+(aq) -> Mg^{2+}(aq) + H2(g)}$$
Why did they ‘split’ the $$\ce{H2+(aq)}$$ ion into $$\ce{2H+(aq)}$$?
• Because there is no such thing as $\ce{H2+(aq)}$. – Ivan Neretin Jan 15 at 20:59
$$\ce{H2+ (aq)}$$ ions don’t exist. Furthermore, even if they did, they would not make a neutral compound $$\ce{H2SO4}$$ with the sulfate ion, since the sulfate ion has a $$2-$$ charge.
So you must split $$\ce{H2SO4(aq)}$$ into $$\ce{2 H+}$$ and $$\ce{SO4^2-}$$ ions.