###Introduction###
In laboratory scenarios, it is common for scientists to add indicators to their solutions. An **indicator** is a special liquid that changes the color of the solution based on the pH of the solution. It is important to note that the indicator itself does not contribute to pH change.
This article will use some actual indicators for use in solving acid-base chemistry problems. These problems will for the most part involve using the color of the indicator to estimate the pH of the solution, or vice-versa.
###Common Indicators###
There are two types of indicators that are particularly common in the laboratory. The first of these is known as the **universal indicator** for its extremely frequent use. The color of the indicator ranges from red for very acidic solutions to purple for very basic solutions. The full color table is available here: http://www.docbrown.info/page03/AcidsBasesSalts/pHscale.gif.
The second very common indicator is known as **phenolphthalein**. This indicator works very differently from the universal indicator in that there are only two possible colors for the indicator: colorless and pink. As the solution gets increasingly basic, the indicator will turn the solution a darker pink. The most acidic of solutions have very little to no color added to them by the indicator.
Example 1: You are working with an aqueous solution of an ionic compound that is naturally pink in color. Which indicator should you use to indicate changes in pH? Explain.
Solution: The actual color of the solution is the same as the color in the phenolphthalein indicator. Even though the “pinkness” of the solution may vary a bit, usage of the universal indicator will yield more apparent changes in color. Thus you should use the universal indicator.
###Using Indicator Color to Estimate pH###
This is the main objective of using indicators in the first place, perhaps when pH meters are not available or simply to serve as a “check” that your sensor is working.
Example 2: You grab a bottle labeled “universal indicator” and put some drops into an aqueous solution of sulfuric acid \(H_2SO_4\). The pH meter reads a pH of 13, while the normally colorless solution turns dark pink in presence of the indicator. This is not what you expected to happen! Give two possible explanations as to why you are getting such bizarre results.
Solution: If the pH was 13, then a dark pink solution would make perfect sense if you were using the phenolphthalein indicator. Thus it is possible that the indicator was incorrectly labeled. Also, it is possible that the pH meter is broken and is giving faulty readings—extremely acidic solutions give a reddish color when tinted with universal indicator.
Here is a different kind of question that depicts a different laboratory scenario.
Example 3: You have a solution of aqueous sodium hydroxide with an unknown pH. You cannot measure the pH with a pH meter because the one you have access to is broken. To get an idea of an estimate for the pH, you add a few drops of universal indicator. The normally colorless solution turns a light purple. Estimate the pH of the solution to the nearest whole number.
Solution: Well, we can refer to the universal indicator color chart. We see that a light purple coloration corresponds to a pH of about 12.
Example 4: A solution tinted with phenolphthalein that is colorless without the indicator remains colorless. Which of the following is the most likely pH of the solution?
A. 1
B. 7
C. 13
Solution: Since the solution is colorless with the phenolphthalein indicator, it is acidic according to the color chart for this indicator. Thus a pH of 1 is the most likely and realistic.
###Using pH to Predict Indicator Color###
This is the reverse procedure of what was shown in the previous section. If you are using the pH to predict the indicator’s color, it is essential to know what indicator you are using.
Example 5: You add universal indicator to a solution with a pH of 9. What will the color of the solution be if it is colorless without any indicator?
Solution: According to the Universal Indicator color chart, the otherwise colorless solutions with pH values of 9 tend to be dark green. Simply refer to what the table says here.
Example 6: You add an unknown indicator to a colorless solution with a pH of 7. The solution turns a light blue. Using this table of indicator colors, identify which indicator you are most likely using. The needed table is available at http://www.writework.com/uploads/12/127232/image26.png.
Solution: On the table, view pH = 7 and scan the column to find an indicator that turns blue for this pH value. The litmus indicator is the most likely since it is the only one that is blue at pH = 7.
###Problem-Solving Using Indicators###
This section of the article will combine the concept of indicators with other concepts in acid-base chemistry to create more complicated problems. Work through them carefully, as they involve quite a bit of algebra.
Example 7: A solution of sodium hydroxide (a strong base) has a concentration of hydroxide ions of \(1.5 \times 10^{-3}\) molars. The solution is split between two containers.
Solution: To identify the color of the solution as some indicator is added, we need the pH. Recall that
$$pH = -\log_{10}[H^{+}]$$
We are given that $[OH^{-}] = 1.5 \times 10^{-3}$ molars. Furthermore, we know that \\([H^{+}][OH^{-}] = 1.0 \times 10^{-14}\\), so with this information, we can find the pH.
$$[H^{+}] = \frac{1.0 \times 10^{-14}}{[OH^{-}]} = \frac{1.0 \times 10^{-14}}{1.5 \times 10^{-3}} = 6.7 \times 10^{-12}$$
$$pH = -\log_{10}[H^{+}] = -\log_{10}(6.7 \times 10^{-12}) = 11.17$$
a. According to the phenolphthalein table, this solution will be pink.
b. According to the Universal Indicator table, a solution with a pH close to 11 will be dark blue in color.
c. As 90% of the solutions have evaporated, 90% of the liquid water in each solution has evaporated, meaning the concentration of base in each solution has increased by a factor of 10. In turn the concentration of acid in each solution decreased by a factor of 10. The new pH of both solutions is
$$pH = -\log_{10}[H^{+}] = -\log_{10}(6.7 \times 10^{-13}) = 12.17$$
The colors of the solutions will change accordingly. According to the respective tables, the solution with universal indicator will turn purple. However, the solution with the phenolphthalein will become an even darker pink as it becomes more basic.
Example 8: You are performing an acid-base titration between .400 liters of \(.08\) molar magnesium hydroxide and 1.000 liters of \(.05\) molar hydroiodic acid. The initial basic solution has a few drops of universal indicator added to it.
Solution:
a. The pH of the initial solution can be found with the concentration of the initial basic solution.
$$pOH = -\log_{10}[OH^{-}] = -\log_{10}(.08) = 1.10$$
Then
$$pH = 14 - pOH = 12.90$$
According to the table, a solution with a pH near $13$ will be purple.
b. The reaction is represented by the equation
$$Mg(OH)_2 + HCl \rightarrow MgCl_2 + H_2O$$
Balancing this equation gives
$$Mg(OH)_2 + 2HCl \rightarrow MgCl_2 + 2H_2O$$
We must now find the number of moles of acid and base that are used in total.
Acid:
$$.400 L \times \frac{.08 \; mol}{1 \; L} = .03200 \; mol$$
Base:
$$1.000 L \times \frac{.05 \; mol}{1 \; L} = .05 \; mol$$
That means, when the reaction is complete, there will be \\(.05 - .03200 = .018 \; mol\\) of base in excess. The total volume of the solution is 1.400 liters. Therefore we can calculate the new molarity [of hydroxide ions]:
$$[OH^{-}] = \frac{.018 \; mol}{1.400 \; L} = .013 \; M$$
Now we can find the new pOH:
$$pOH = -\log_{10}[OH^{-}] = -\log_{10}(0.13 \; M) = 1.89$$
Thus the new pH is
$$14 - pOH = 12.11$$
As the pH is near \(12\), the solution will be purple.
c. A green solution tinted with universal indicator, according to the table, means the pH of the solution is \(7\), i.e. neutral. Therefore the number of moles of acid and base must be the same for full neutralization, requiring \(.05\) moles of acid. This is infeasible for the given setup because we only have \(.03200\) moles of acid.
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Indicators are a very powerful tool in laboratory settings. They give another dimension to understanding pH, and have very practical uses in the laboratory.
###Sources###
DocBrown.edu. Web, accessed Universal Indicator, http://www.docbrown.info/page03/AcidsBasesSalts/pHscale.gif
Marx, Graig. Personal interview. May 20, 2014.
Mindset. Web, accessed May 27, 2014, http://www.mindset.co.za/assets/SCIUNIT21_TB01.gif
Writework.com. Seven different indicators: http://www.writework.com/uploads/12/127232/image26.png