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

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Answer 1 · 2017-07-12T18:06:42

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

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?