which one of the following cannot disproportionate cl-, clo3-, cl2, cu+, fe2+
In order to disproportionate the ion must be in the "middle" of its range; i.e., Cl2 is at zero. It can go down (to -1) or up to +1 or higher. Fe^2+ can go down (to +1 or 0) or up (to +3). So the one you're looking for is either at the top or the bottom of its range.
To determine which of the following species cannot undergo disproportionation, we need to understand what disproportionation means.
Disproportionation is a redox reaction where an element in a particular oxidation state is simultaneously reduced and oxidized, resulting in the formation of two different oxidation states of the same element.
Let's analyze each species to see if they can undergo disproportionation:
1. Cl- (chloride ion): Chloride ions have an oxidation state of -1. Since they do not contain chlorine in two different oxidation states, they cannot undergo disproportionation.
2. ClO3- (chlorate ion): Chlorate ions have an oxidation state of +5. These ions can undergo disproportionation and are commonly used in laboratory reactions.
3. Cl2 (chlorine gas): Chlorine gas has an oxidation state of 0. It can undergo disproportionation, as seen in reactions such as the reaction with water to produce HCl and HClO.
4. Cu+ (copper(I) ion): Copper(I) ions have an oxidation state of +1. They do not contain copper in two different oxidation states, so they cannot undergo disproportionation.
5. Fe2+ (iron(II) ion): Iron(II) ions have an oxidation state of +2. They can also undergo disproportionation reactions, forming Fe3+ and Fe0 species.
From the analysis above, we can conclude that the species Cl- (chloride ion) and Cu+ (copper(I) ion) cannot undergo disproportionation, as they do not contain the respective elements in two different oxidation states.