To solve for pH, we use the **Henderson-Hasselbalch equation**:

$\overline{){\mathbf{pH}}{\mathbf{=}}{{\mathbf{pK}}}_{{\mathbf{a}}}{\mathbf{+}}{\mathbf{log}}\mathbf{\left(}\frac{\mathbf{conjugate}\mathbf{}\mathbf{base}}{\mathbf{weak}\mathbf{}\mathbf{acid}}\mathbf{\right)}}$

We are given the following:

**pK _{a} = 3.00In^{-} = 7.00%HIn = 93.00%**

Since in the equation, we just need to compare the ratio of conjugate and weak acid present, we can directly plug-in the given percentages to the equation, thus we **solve for pH**:

$\overline{){\mathbf{pH}}{\mathbf{=}}{{\mathbf{pK}}}_{{\mathbf{a}}}{\mathbf{+}}{\mathbf{log}}\mathbf{\left(}\frac{\mathbf{conjugate}\mathbf{}\mathbf{base}}{\mathbf{weak}\mathbf{}\mathbf{acid}}\mathbf{\right)}}\phantom{\rule{0ex}{0ex}}\mathbf{pH}\mathbf{=}\mathbf{3}\mathbf{.}\mathbf{00}\mathbf{+}\mathbf{log}\mathbf{\left(}\frac{\mathbf{7}\mathbf{.}\mathbf{00}\overline{)\mathbf{\%}}}{\mathbf{93}\mathbf{.}\mathbf{00}\overline{)\mathbf{\%}}}\mathbf{\right)}\phantom{\rule{0ex}{0ex}}\mathbf{pH}\mathbf{=}\mathbf{3}\mathbf{.}\mathbf{00}\mathbf{+}\mathbf{(}\mathbf{-}\mathbf{1}\mathbf{.}\mathbf{123}\mathbf{)}\phantom{\rule{0ex}{0ex}}\mathbf{pH}\mathbf{=}\mathbf{3}\mathbf{.}\mathbf{00}\mathbf{-}\mathbf{1}\mathbf{.}\mathbf{123}$

A certain indicator HIn has a pK_{a} of 3.00 and a color change becomes visible when 7.00% of the indicator has been converted to In^{-}. At what pH is this color change visible?

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Our tutors have indicated that to solve this problem you will need to apply the Acid Base Indicators concept. If you need more Acid Base Indicators practice, you can also practice Acid Base Indicators practice problems.

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Our data indicates that this problem or a close variation was asked in Chemistry: An Atoms First Approach - Zumdahl Atoms 1st 2nd Edition. You can also practice Chemistry: An Atoms First Approach - Zumdahl Atoms 1st 2nd Edition practice problems.