The reason why a pH indicator changes color depending on the acidity (or alkalinity) of the solution is due to the chemical structure of the indicator molecule and its response to changes in hydrogen ion (H+) concentration.
pH indicators are typically weak acids or bases that undergo a reversible chemical reaction between their acidic and basic forms. The indicator molecule exists in an equilibrium between its protonated form (HIn) and its deprotonated form (In-), depending on the presence of hydrogen ions in the solution. The equilibrium can be represented as follows:
HIn ⇌ In- + H+
In the presence of acidic conditions (high hydrogen ion concentration), the equilibrium is shifted towards the protonated form (HIn). This acidic form often has one color. As the solution becomes less acidic and more alkaline (low hydrogen ion concentration), the equilibrium shifts towards the deprotonated form (In-), which typically has a different color. This color change is what we observe as the pH indicator "changing color."
The specific color change that occurs depends on the chemical structure of the indicator molecule. Different indicators have different pH ranges over which they change color, known as their transition range. For example, phenolphthalein is commonly used as an indicator in the pH range of about 8.2 to 10.0. It appears colorless in acidic solutions and pink in alkaline solutions within that range.
The color change is a result of the molecular structure of the indicator being influenced by the presence or absence of hydrogen ions. The presence of hydrogen ions leads to a different electronic distribution within the molecule, altering its absorption and reflection of light, and thus resulting in a different observed color.
It's important to note that pH indicators are typically used for approximate measurements of pH and are not as precise as pH meters or pH electrodes. However, they provide a convenient visual indication of the approximate acidity or alkalinity of a solution based on the observed color change.