Help please?
can the following reaction be classified as a redox reaction (1)? what factors did you look at to make this determination? is anything being reduced/ oxidized? if so, what?
cl2(g) + 2nabr(aq) > 2nacl(aq) + br2(aq)
Cl₂ + 2eˉ => 2Clˉ < = > Reduction (gain of 2 electrons)
2Brˉ => Br₂ + 2eˉ < = > Oxidation (loss of 2 electrons)
Remember OIL RIG …
Oxidation Is Loss (of eˉ) Reduction Is Gain (of eˉ)
Sodium (Na) is a spectator ion; i.e., doesn't undergo redox.
To determine whether the given reaction is a redox reaction, you need to analyze the changes in the oxidation states of the elements involved. Here's how you can do it:
1. Assign oxidation states:
Determine the oxidation state (or oxidation number) of each element in the reaction. Oxidation states reflect the apparent charge that each atom would have if electrons were completely transferred in the bond.
For this reaction, chlorine (Cl) is a diatomic molecule and is assigned an oxidation state of 0 in its elemental form. Sodium (Na) is in Group 1 of the periodic table and has an oxidation state of +1. Bromine (Br) also has an oxidation state of 0 in its elemental form. Oxygen (O) is usually assigned an oxidation state of -2, and hydrogen (H) has an oxidation state of +1 when combined with non-metals.
2. Determine changes in oxidation states:
Compare the oxidation states of each element in the reactants and products. If the oxidation state of an element increases, it is being oxidized, and if it decreases, it is being reduced.
In the given reaction:
- The oxidation state of chlorine changes from 0 to -1 in NaCl, so it is being reduced (gaining electrons).
- For bromine, its oxidation state changes from 0 to -1 in NaBr, so it is being reduced as well.
- Sodium has an oxidation state of +1 in both NaBr and NaCl, so it is not being oxidized or reduced.
Based on these observations, we can conclude that the given reaction is a redox reaction. Chlorine and bromine are being reduced, while sodium remains unchanged in terms of its oxidation state.