To determine what most likely happens during the given chemical reaction:
[tex]\[
Mg_{(s)} + H^+ Cl^- \rightarrow Mg^{2+} Cl^- + H_2
\][/tex]
We need to analyze the changes in the oxidation states of the elements involved.
1. Magnesium (Mg):
- In its elemental form, magnesium (Mg) is in the zero oxidation state ([tex]\(Mg_{(s)}\)[/tex]).
- After the reaction, magnesium is present as [tex]\(Mg^{2+}\)[/tex], indicating it is in the +2 oxidation state.
Since magnesium goes from 0 to +2, it means that it has lost two electrons. The change can be written as:
[tex]\[
Mg \rightarrow Mg^{2+} + 2e^-
\][/tex]
This illustrates that magnesium loses two electrons during the reaction.
2. Hydrogen (H):
- The hydrogen ion ([tex]\(H^+\)[/tex]) starts in the +1 oxidation state.
- In the product, hydrogen is present as [tex]\(H_2\)[/tex] gas, which is in the 0 oxidation state because [tex]\(H_2\)[/tex] is a diatomic molecule with no charge.
The reduction of hydrogen can be described as:
[tex]\[
2H^+ + 2e^- \rightarrow H_2
\][/tex]
Hydrogen gains two electrons in this case, but the problem specifically concerns what happens to magnesium and chlorine.
3. Chlorine (Cl):
- Throughout the reaction, chlorine starts and remains as [tex]\(Cl^-\)[/tex] with a -1 oxidation state.
There is no change in the oxidation state of chlorine, so it neither gains nor loses electrons.
Summarizing from the steps:
- Magnesium loses two electrons.
- Hydrogen gains electrons.
- Chlorine’s oxidation state does not change.
Hence, based on the reaction, the correct and most likely event is:
[tex]\[
\text{Magnesium loses two electrons.}
\][/tex]
