In the reaction Bromine
`3Br_2 + 6 OH^(-) rarr 5 Br ^(-) + BrO_3^(-) + 3H_2O`

Options

A. is reduced
B. is oxidized
C. disproportionates
D. disintegrates

Answer

To determine whether bromine is reduced, oxidized, disproportionated, or disintegrated in the given reaction, we need to analyze the oxidation states of bromine in the reactants and products.

1. **Identify the oxidation state of bromine in the reactants**:
- In the reactants, bromine (Br₂) is in its elemental form, which means it has an oxidation state of 0.

2. **Identify the oxidation states of bromine in the products**:
- In the products, we have two different forms of bromine:
- **Br⁻**: The oxidation state of bromine here is -1 (since it has a negative charge).
- **BrO₃⁻**: To find the oxidation state of bromine in this compound, we can set up the equation:
- Let the oxidation state of bromine be \( x \).
- The oxidation states of oxygen are -2. Since there are three oxygen atoms, the total contribution from oxygen is \( 3 \times (-2) = -6 \).
- The overall charge of the bromate ion (BrO₃⁻) is -1.
- Therefore, we can write the equation:
\[ x - 6 = -1 \]
- Solving for \( x \):
\[ x = -1 + 6 = +5 \]
- Thus, the oxidation state of bromine in BrO₃⁻ is +5.

3. **Determine the changes in oxidation states**:
- Bromine changes from an oxidation state of 0 in Br₂ to -1 in Br⁻ (reduction).
- Bromine also changes from an oxidation state of 0 in Br₂ to +5 in BrO₃⁻ (oxidation).

4. **Conclusion**:
- Since bromine undergoes both reduction (from 0 to -1) and oxidation (from 0 to +5), it is undergoing a disproportionation reaction. In a disproportionation reaction, a single element is both oxidized and reduced.

Thus, the correct answer is that bromine **disproportionates** in the given reaction.