Are you interested in learning more about titration flask? Contact us today to secure an expert consultation!
Indicators are substances whose solutions change color due to changes in pH. These are called acid-base indicators. They are usually weak acids or bases, but their conjugate base or acid forms have different colors due to differences in their absorption spectra. Did you know that the color of hydrangea flowers depends on the pH of the soil in which they are grown?
Figur
e \(\PageIndex{1}\): This picture shows various colors of hydrangea flowers. (CC BY-SA 4.0 International;
FlowerPowerH2020 via Wikipedia).Indicators are organic weak acids or bases with complicated structures. For simplicity, we represent a general indicator by the formula \(\mathrm{\color{Blue} HIn}\), and its ionization in a solution by the equilibrium,
\[\mathrm{ {\color{Blue} HIn} \rightleftharpoons H^+ + {\color{Red} In^-}}\]
and define the equilibrium constant as Kai,
\[K_{\large\textrm{ai}} = \mathrm{\dfrac{[H^+][{\color{Red} In^-}]}{[{\color{Blue} HIn}]}}\]
which can be rearranged to give
\[\mathrm{\dfrac{[{\color{Red} In^-}]}{[{\color{Blue} HIn}]}} = \dfrac{K_{\large\textrm{ai}}}{\ce{[H+]}}\]
When \(\ce{[H+]}\) is greater than 10 Kai, \(\mathrm{\color{Red} In^-}\) color dominates, whereas color due to \(\mathrm{\color{Blue} HIn}\) dominates if \(\ce{[H+]} < \dfrac{K_{\large\textrm{ai}}}{10}\). The above equation indicates that the color change is the most sensitive when \(\ce{[H+]} = K_{\large\textrm{ai}}\) in numerical value.
We define pKai = - log(Kai), and the pKai value is also the pH value at which the color of the indicator is most sensitive to pH changes.
Taking the negative log of Kai gives,
\[-\log K_{\large\textrm{ai}} = -\log\ce{[H+]} - \log\mathrm{\dfrac{[{\color{Red} In^- }]}{[{\color{Blue} HIn}]}}\]
or
\[\mathrm{pH = p\mathit K_{\large{ai}}} + \log\mathrm{\dfrac{[{\color{Red} In^-}]}{[{\color{Blue} HIn}]}}\]
This is a very important formula, and its derivation is very simple. Start from the definition of the equilibrium constant K; you can easily derive it. Note that pH = pKai when \([\mathrm{\color{Red} In^-}] = [\mathrm{\color{Blue} HIn}]\). In other words, when the pH is the same as pKai, there are equal amounts of acid and base forms. When the two forms have equal concentration, the color change is most noticeable.
Colors of substances make the world a wonderful place. Because of the colors and structures, flowers, plants, animals, and minerals show their unique characters. Many indicators are extracted from plants. For example, red cabbage juice and tea pigments show different colors when the pH is different. The color of tea darkens in a basic solution, but the color becomes lighter when lemon juice is added. Red cabbage juice turns blue in a basic solution, but it shows a distinct red color in an acidic solution.
Common Indicators: Some common indicators and their pKai (also referred to as pKa) values are given in a table form. Name Acid Color pH Range of Color Change Base Color Methyl violet Yellow 0.0 - 1.6 Blue Thymol blue Red 1.2 - 2.8 Yellow Methyl orange Red 3.2 - 4.4 Yellow Bromocresol green Yellow 3.8 - 5.4 Blue Methyl red Red 4.8 - 6.0 Yellow Litmus Red 5.0 - 8.0 Blue Bromothymol blue Yellow 6.0 - 7.6 Blue Thymol blue Yellow 8.0 - 9.6 Blue Phenolphthalein Colorless 8.2 - 10.0 Pink Thymolphthalein Colorless 9.4 - 10.6 Blue Alizarin yellow R Yellow 10.1 - 12.0 RedThere is a separate file for this, and it can also be accessed from the Chemical Handbook menu.
Find an indicator for the titration of a 0.100 M solution of a weak acid \(\ce{HA}\) (\(K_a = 6.2 \times 10^{-6}\)) with 0.100 M \(\ce{NaOH}\) solution.
First, you should estimate the pH at the equivalence point, at which the solution is 0.0500 M \(\ce{NaA}\). This is a hydrolysis problem, but the following method employs the general principle of equilibrium.
\[\begin{array}{ccccccc}
\ce{A- &+ &H2O &\rightleftharpoons &HA &+ &OH-}\\
0.0500-y &&&&y &&y
\end{array} \nonumber\]
If we multiply the numerator and the denominator by \(\ce{[H+]}\), rearrange the terms, note that \(\ce{[H+][OH- ]} = K_{\large\textrm w}\), and by the definition of Ka of the acid, we have the following relationship:
\[\dfrac{y^2}{0.0500-y} = \ce{\dfrac{[HA][OH- ]}{[A- ]} \dfrac{[H+]}{[H+]}} = \dfrac{K_{\large\textrm w}}{K_{\large\textrm a}}\]
\[\begin{align*}
y &= \left(0.0500\left(\dfrac{K_{\large\textrm w}}{K_{\large\textrm a}}\right)\right)^{1/2}\\
&= 9.0 \times 10^{-6}
\end{align*}\]
\[\begin{align*}
\ce{pOH} &= -\log \ce{[OH- ]} = -\log 9.0 \times 10^{-6}\\
&= 5.05\\ \\
\ce{pH} &= 14 - 5.05\\
&= 8.95
\end{align*}\]
Phenolphthalein in the table above has a pKai value of 9.7, which is the closest to the pH of equivalence point in this titration. This indicator is colorless in acidic solution, but a light PINK appears when the pH is > 8. The color becomes more INTENSE PINK as the pH rises. A parade of the color intensities is shown below:
___The equivalence point is when the color changes most rapidly, not when the solution has changed color. Improper use of indicators will introduce inaccuracy to titration results.
Want more information on elisa plates? Feel free to contact us.
Indicators change color gradually at various pH. Let us assume that the acid form has a blue color and the basic form has red color. The variation of colors at different pH is shown below. The background color affects their appearance and our perception of them.
RnB RnB RnB RnB RnB RnB RnB RnB RnB RnB RnB RnB RnB RnB RnB RnB RnB RnB RnB RnB RnB RnB RnB RnB RnB RnB RnB RnB RnB RnB _______The long stretched color in the middle of the last line has equal intensity of BLUE and RED. If a solution has a color matching this, the pH would be the same as the pKai of the indicator, provided that the conjugate forms of the indicator have the BLUE and RED colors.
Chung (Peter) Chieh (Professor Emeritus, Chemistry @ University of Waterloo)
The analyte is prepared by dissolving the substance being studied into a solution. The solution is usually placed in a flask for titration. A small amount of indicator is then added into the flask along with the analyte. The reagent is usually placed in a burette and slowly added to the analyte and indicator mixture. The amount of reagent used is recorded when the indicator causes a change in the color of the solution.
Some titrations requires the solution to be boiled due to the \(CO_2\) created from the acid-base reaction. The \(CO_2\) forms carbonic acid (\(H_2CO_3\)) when dissolved in water that then acts as a buffer, reducing the accuracy of data. After boiling water, most of the \(CO_2\) will be removed from the solution allowing the solution to be titrated to a more accurate endpoint. The endpoint is the point where all of the analyte has be reacted with the reagent.
A useful indicator has a strong color that changes quickly near its pKa. These traits are desirable so only a small amount of an indicator is needed. If a large amount of indicator is used, the indicator will effect the final pH, lowering the accuracy of the experiment. The indicator should also have a pKa value near the pH of the titration's endpoint. For example a analyte that is a weak base would require an indicator with a pKa less than 7. Choosing an indicator with a pKa near the endpoint's pH will also reduce error because the color change occurs sharply during the endpoint where the pH spikes, giving a more precise endpoint.
Figure 1: A Basic Titration Curve, The horizontal lines show the range of pH in which phenolphthalein (blue) and methyl orange (red) changes color. The middle line represents the pKa, while the two outer lines represent the end or start of the color changes. The peak and light blue highlights show the range in which the color changes will occur based on the amount of titrant added.
Notice that this reaction is between a weak acid and a strong base so phenolphthalein with a pKa of 9.1 would be a better choice than methyl orange with a pKa of 3.8. If in this reaction we were to use methyl orange as the indicator color changes would occur all throughout the region highlighted in pink. The data obtained would be hard to determine due to the large range of color change, and inaccurate as the color change does not even lie with the endpoint region. Phenolphthalein on the other hand changes color rapidly near the endpoint allowing for more accurate data to be gathered.
Figure 2: Titration Demonstration, The picture was taken during a vinegar titration lab. C2H4O2(aq) - acetic acid- was titrated against NaOH (aq) - sodium hydroxide - using phenolphthalein as indicator. The image on the right is submicroscopic view of the titration reaction featuring C2H4O2(aq), which is partially dissociated and NaOH (aq) that is completely dissociated into Na+(aq) in purple and OH-(aq).
Multiply the volume of reagent added to get to the endpoint, with the molarity of the reagent to find the moles of reagent used. With the balanced equation of the acid-base reaction in question to find the moles of unknown substance. Then the original molarity can be calculated by dividing through with the initial volume.
Note
For example an unknown molarity of \(HCl\) acts as the analyte. 50 mL of it is placed into a flask and a 0.1 M solution of \(NaOH\) will be the reagent. The endpoint is \(pH= 7\) so litmus, with a pKa of 6.5 is chosen. The color of the solution changes when 10 mL of 0.1 M \(NaOH\) is added.
The balanced neutralization reaction:
\[ HCl_{(aq)} + NaOH_{(aq)} \rightarrow H_2O_{(l)} + Na^+ + Cl^-\]
Or just the net ionic equation
\[H^+ + OH^- \rightarrow H_2O_{(l)} \]
The following equation can then be derived
\[ X= (0.1\; M \;\; NaOH) (10\;mL) \left(\dfrac{1\;L}{1000\; mL} \right) \left( \dfrac{1\;mol\; NaOH}{1\; mol \; OH^-} \right)\]
\[H^+ = X \;\; HCl\]
\(X = 0.0010 \;mol \;of\; HCl\)
The molarity is now easily solved for
\[ \dfrac{0.0010\; mol\; HCl}{0.050\; L} = 0.020\;M \;HCl\]
If you are looking for more details, kindly visit t175 flasks.