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:
  $$\ce{Mg(s) + H2SO4(aq) -&gt; MgSO4(aq) + H2(g)}$$


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

This was the answer:


  $$\ce{Mg(s) + 2H+(aq) -&gt; Mg^{2+}(aq) + H2(g)}$$


Why did they ‘split’ the $\ce{H2+(aq)}$ ion into $\ce{2H+(aq)}$?
$\ce{H2+ (aq)}$ 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.