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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}

This was the answer:

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

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

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    $\begingroup$ Because there is no such thing as $\ce{H2+(aq)}$. $\endgroup$ – Ivan Neretin Jan 15 at 20:59
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$\ce{H2+ (aq)}$ 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.

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