Methyl orange, \(\mathrm{HMO}\), is a common acid-base indicator. In solution it
ionizes according to the equation: $$\begin{array}{c}
\mathrm{HMO}(\mathrm{aq})=\mathrm{H}^{+}(\mathrm{aq})+\mathrm{MO}^{-}(\mathrm{aq})
\\\
\mathrm{red}\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad\mathrm{yellow}
\end{array}$$
a. Why does adding \(6 \mathrm{M}\) \(\mathrm{HCI}\) to the yellow solution of
methyl orange tend to cause the color to change to red? (Note that in solution
\(\mathrm{HCl}\) exists as \(\mathrm{H}^{+}\) and \(\mathrm{Cl}^{-}\) ions.)
b. Why does adding \(6 \mathrm{M}\) \(\mathrm{NaOH}\) to the red solution tend to
make it turn back to yellow? Note that in solution \(\mathrm{NaOH}\) exists as
\(\mathrm{Na}^{+}\) and \(\mathrm{OH}^{-}\) ions. (Hint: How does increasing
\(\left[\mathrm{OH}^{-}\right]\) shift Reaction 3 in the discussion section? How
would the resulting change in \(\left[\mathrm{H}^{+}\right]\) affect the
dissociation reaction of \(\mathrm{HMO}\)? Explain these as part of your
answer.)
