Question

How do we use the method of half-equations, and redox formalism to represent the oxidation of magnesium metal by hydrochloric acid?

Answer 1

The `"oxidation state"` of magnesium metal is a big fat `"ZERO"`........ And the metal oxidation state in `MgCl_2` is `+II`.

Explanation:

You have performed a redox reaction, the which we could formally separate into oxidation/reduction half equations.......

`Mg(s) rarr Mg^(2+) +2e^-` `(i)`, i.e. the electron rich metal formally LOSES 2 valence electrons

And a reduction half-equation.......

`HCl(aq) + e^(-) rarr 1/2H_2(g)uarr +Cl^-` `(ii)`

A oxidizing protium ion formally gains one electron........

We add `(i)` and `(ii)` such that electrons DO NOT APPEAR in the final redox equation, i.e. `(i) + 2xx(ii)` gives.........

`Mg(s) + 2HCl(aq) rarr MgCl_2(aq) + H_2(g)uarr`

........alternatively.......

`Mg(s) + 2HCl(aq) rarr Mg^(2+) + 2Cl^(-) + H_2(g)uarr`

Charge and mass are CONSERVED, as is absolutely required for ANY chemical reaction.......

Capisce?