A net cell reaction is given as
`Cr+3H_(2)O+OCl^(-)rarrCr^(3+)+3Cl^(-)+6OH^(-)`
The species undergoing reduction is

Options

A. Cr
B. `Cl^(-)`
C. `OCl^(-)`
D. None of these

Explanation

To determine the species undergoing reduction in the given net cell reaction:

**Given Reaction:**
\[ \text{Cr} + 3\text{H}_2\text{O} + \text{OCl}^- \rightarrow \text{Cr}^{3+} + 3\text{Cl}^- + 6\text{OH}^- \]

**Step 1: Identify the oxidation states of the elements involved.**
- Chromium (Cr) in the reactants is in the elemental state, so its oxidation state is 0.
- In the products, chromium is in the form of \(\text{Cr}^{3+}\), which has an oxidation state of +3.
- For \(\text{OCl}^-\), we need to determine the oxidation state of chlorine (Cl). The oxidation state of oxygen (O) is -2, and the overall charge of \(\text{OCl}^-\) is -1. Therefore, we can set up the equation:
\[ x + (-2) = -1 \implies x = +1 \]
Thus, the oxidation state of Cl in \(\text{OCl}^-\) is +1.

**Step 2: Analyze the changes in oxidation states.**
- Chromium goes from 0 to +3, indicating that it is being oxidized (loss of electrons).
- Chlorine in \(\text{OCl}^-\) goes from +1 to -1 (in \(\text{Cl}^-\)), indicating a decrease in oxidation state, which means it is being reduced (gain of electrons).

**Step 3: Identify the species undergoing reduction.**
- Since the oxidation state of chlorine decreases from +1 in \(\text{OCl}^-\) to -1 in \(\text{Cl}^-\), we conclude that \(\text{OCl}^-\) is the species that undergoes reduction.

**Final Answer:**
The species undergoing reduction is \(\text{OCl}^-\).

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