When we add NaOH in the first trial, this will dissolve into ions, including OH ^- , which will combine with the existing H ⁺ ions to make water. According to Le Chatelier's principle, the reaction's equilibrium will oppose the reduction in the concentration of products, temporarily increasing the rate of the forward reaction. The newly established equilibrium is then shifted to the right, i.e. in favour of the products.

By adding HCl in the second trial we increase the concentration of H ⁺ ions, which will combine with the OH ^- ions to make water. Same as with the NaOH , the reaction's equilibrium will shift right to compensate for the change in concentration.

According to our observations, the NaCl trial turned orange, which would suggest an equilibrium shift to the reactants. However, I believe that we conducted this trial incorrectly (perhaps by adding the wrong substance), since I cannot identify any chemical mechanism for this shift. In particular, Na ⁺ and Cl ^- do not seem to react with any substances in the equation above, and do not form any kind of precipitate. Therefore, adding NaCl should not have had any effect on the equilibrium above, with the colour remaining green throughout (). This also makes sense because BTB is a an acid-base indicator and NaCl is neutral; therefore we would not expect a colour change since the pH does not change.