The reasons for this are explored more thoroughly in Unit 4 Equilibrium. Although you won't explicitly learn this in Unit 4, the theory will help to explain the concept better.

The basic idea is that there is a reaction:

• HI (indicator in acidic form; coloured) ↔ H⁺ + I^- (indicator in basic form; coloured a different colour)

Basically, the HI and I^- compounds are conjugate acid/bases, and they are coloured, which helps to serve their purpose as indicators. Also note that one of HI or I^- can also be colourless, as only one needs to be coloured for it to serve its purpose.

Tip: Note here that HI and I^- are actually acids/bases, which is why you're usually told to only add a few drops - otherwise you will actually affect the acid-base system you're trying to observe too much - the indicator itself is an acid or a base.

However, HI and I^- will co-exist but at different percentages depending on the "position of the reaction" (may make no sense without knowledge of equilibrium).

• For example, if it's very acidic, then the reaction will be driven to the left (since there is lots of H⁺ driving the backwards reaction), which will create a lot of HI.

• However, if it were slightly less acidic, but still acidic, there might be a lot of HI, but maybe a little bit of I^- occurring.

This means there will be two coloured compounds co-existing, which will cause ambiguity (i.e.: is it red or yellow?). For that reason, most likely previous chemists have defined a particular pH range when it is "officially" yellow or "officially" red. (This range would probably be based on some calculated percentage that one of the compounds must reach before it becomes "officially" one colour or the other)

This answer is just for curiosity's sake, and my interpretation of how indicators work. I can say with 99.9999% confidence that you wouldn't be examined on this.