Let's analyze the given chemical reaction to determine which statement is true:
[tex]\[ \text{H}_2 + \text{O}_2 \longrightarrow \text{H}^+ + \text{O}^{2-} \][/tex]
1. Balancing Atoms:
- Reactants: On the left side, we have:
[tex]\[ 2 \text{H atoms (from H}_2\text{)} \][/tex]
[tex]\[ 2 \text{O atoms (from O}_2\text{)} \][/tex]
- Products: On the right side, we have:
[tex]\[ 1 \text{H atom (from H}^+\text{)} \][/tex]
[tex]\[ 1 \text{O atom (from O}^{2-}\text{)} \][/tex]
Clearly, the number of atoms on the reactant side does not match the number of atoms on the product side. We have 2 hydrogen atoms and 2 oxygen atoms on the left, but only 1 hydrogen atom and 1 oxygen atom on the right. Thus, the reaction is not balanced in terms of the number of atoms.
2. Oxidation States:
- Reactants:
- For [tex]\(\text{H}_2\)[/tex]: The oxidation state of H in [tex]\(\text{H}_2\)[/tex] is 0.
- For [tex]\(\text{O}_2\)[/tex]: The oxidation state of O in [tex]\(\text{O}_2\)[/tex] is 0.
- Products:
- For [tex]\(\text{H}^+\)[/tex]: The oxidation state of H in [tex]\(\text{H}^+\)[/tex] is [tex]\(+1\)[/tex].
- For [tex]\(\text{O}^{2-}\)[/tex]: The oxidation state of O in [tex]\(\text{O}^{2-}\)[/tex] is [tex]\(-2\)[/tex].
Let's consider the changes in oxidation states:
- Hydrogen changes from 0 (in [tex]\(\text{H}_2\)[/tex]) to [tex]\(+1\)[/tex] (in [tex]\(\text{H}^+\)[/tex]), so it is oxidized.
- Oxygen changes from 0 (in [tex]\(\text{O}_2\)[/tex]) to [tex]\(-2\)[/tex] (in [tex]\(\text{O}^{2-}\)[/tex]), so it is reduced.
Since we can also see that the total change in oxidation states isn't balanced (as we need to consider the number of electrons lost and gained should be equal for balancing redox reactions), we can infer that the reaction is also not balanced for oxidation states.
Therefore, the correct statement about the reaction is:
It is not balanced for oxidation state or for number of atoms.
