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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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 $oxidation state$ $"oxidation state"$ of magnesium metal is a big fat $ZERO$ $"ZERO"$ ........ And the metal oxidation state in $M g C l_{2}$ $MgCl_2$ is $+ I I$ $+II$ .

Explanation:

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

$M g (s) \to M g^{2 +} + 2 e^{-}$ $Mg(s) rarr Mg^(2+) +2e^-$ $(i)$ $(i)$ , i.e. the electron rich metal formally LOSES 2 valence electrons

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

$H C l (a q) + e^{-} \to \frac{1}{2} H_{2} (g) ↑ ⏐ ⏐ ⏐ + C l^{-}$ $HCl(aq) + e^(-) rarr 1/2H_2(g)uarr +Cl^-$ $(i i)$ $(ii)$

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

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

$M g (s) + 2 H C l (a q) \to M g C l_{2} (a q) + H_{2} (g) ↑ ⏐$ $Mg(s) + 2HCl(aq) rarr MgCl_2(aq) + H_2(g)uarr$

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

$M g (s) + 2 H C l (a q) \to M g^{2 +} + 2 C l^{-} + H_{2} (g) ↑ ⏐$ $Mg(s) + 2HCl(aq) rarr Mg^(2+) + 2Cl^(-) + H_2(g)uarr$

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

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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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