Natural Indicator
Natural Indicator
Natural Indicator
or plastic wrap and allow the extraction to take place overnight. On the day of lab,
filter or decant the mixture and use the resulting indicator solution in Parts A and B. Alternatively, extraction of the natural
indicators with hot water may be completed as a homework assignment.
Flowers and fruits may be extracted with hot water or 70% isopropyl (rubbing) alcohol at room temperature. There is no
difference in the color charts of natural indicator solutions obtained using these different solvents. Use isopropyl alcohol
only in a well-ventilated lab setting. The recommended minimum extraction time for isopropyl alcohol is 15 minutes
regardless of the solvent used.
Although the extraction step may be the most enjoyable part of this lab for students, it does tend to get a little messy, espe-
cially if the experiment is performed by multiple lab sections in the same room on the same day. Teachers may find it con-
venient to prepare the indicators ahead of time.
A wide range of fruits and flowers contain natural acidbase indicators. The table on the following page summarizes the
information obtained from a brief literature survey. The list is not meant to be exhaustive, but rather to demonstrate the
variety of options suitable for classroom study.
Natural Fruit and Natural Flower Do Not Give Natural
Vegetable indicators Indicators Indicators
Apple skin (red) Dahlias Daffodils
Beets Daylilies Daisies
Blueberries Geraniums Dandelions
Cabbage (red) Hibiscus Marigolds
Cherries Hollyhocks Mums (yellow)
Cranberries Hydrangeas
Grapes (red or purple) Iris (blue)
Onions (red) Morning Glories
Peaches Mums (purple)
Plums Pansies
Radish skin Peonies
Rhubarb skin Petunias
Strawberries Poppies
Tomato leaves Roses (red, pink)
Turnip skin Violets
4
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This experiment reinforces key concepts and definitions in the chemistry of acids and bases. The behavior of natural indi-
cators illustrates the definition of Brnsted acids (proton donors). Most natural indicators are further classified as weak
acids (dissociate only partially in water and their reactions with water are reversible). The different colors observed for
natural indicators thus reflect the position of equilibrium under different conditions. The color transitions are examples
of LeChteliers Principle in action. Use Equation 1 to predict the direction the indicator equilibrium will be shifted as a
result of increasing or decreasing the H
3
O
+
concentration.
A picture is worth a thousand words! Have students draw color charts of indicator color changes using colored pencils.
Discussion
Indicators are dyes or pigments that are isolated from a variety of sources, including plants, fungi, and algae. For example,
almost any flower that is red, blue, or purple in color contains a class of organic pigments called anthocyanins that change color
with pH. The use of natural dyes as acidbase indicators was first reported in 1664 by Sir Robert Boyle in his collection of essays
Experimental History of Colours. Indeed, Boyle made an important contribution to the early theory of acids and bases by using
indicators for the classification of these substances. The idea, however, may actually have originated much earliermedieval
painters used natural dyes treated with vinegar and limewater to make different color watercolor paints.
Acidbase indicators are large organic molecules that behave as weak acidsthey can donate hydrogen ions to water mol-
ecules to form their conjugate bases (Equation 1). The distinguishing characteristic of indicators is that the acid (HIn) and conju-
gate base (In
In
(aq) + H
3
O
+
(aq) Equation 1
(Color A) (Color B)
The abbreviation HIn represents an uncharged indicator molecule, and In
. (3) The solution contains roughly equal amounts of the two forms and the resulting color is intermediate between that
of HIn and In
).
a. What is the pH range in which the most basic form predominates?
b. Calculate the highest H
3
O
+
concentration at which the indicator still exists in this form.
For the red rose indicator as an example, the most basic form of the indicator is yellow in color. This basic form of the
indicator appears to be the predominant form of the indicator in solution at pH values greater than or equal to 10. The
highest H
3
O
+
concentration at which the basic form predominates is therefore 1 10
10
M.
3. For one of the unknown acidbase solutions that you tested, explain why you chose the combination of indicators you did
to determine the pH value of the solution. What is the advantage of using multiple indicators, rather than a single indicator,
to determine the pH of a substance?
In the case of Unknown C (sodium bicarbonate), the natural red rose indicator gave a pH estimate of 89. Bromthymol
blue was selected as an alternate indicator to confirm or narrow the pH range. Bromthymol blue was green, suggesting
that the pH was in the 78 range. The overlap of these two pH estimates suggests that the pH of the solution is very close
to 8. Using multiple indicators, rather than a single indicator, often makes it possible to obtain a more precise (narrow)
estimate of the pH of a substance.
4. Construct a Results Table to summarize the properties of the unknowns.
a. Estimate the pH value of each unknown.
b. Classify each solution as acidic or basic.
c. Within each class of unknownsacids and basesarrange the solutions in order from least acidic to most acidic and
least basic to most basic, respectively.
Analysis of Household Substances
The following household substances provide convenient unknowns for classifying acidbase solutions using indicators.
Alka-Seltzer
Grapefruit juice
Ammonia, household Hair spray
Antacid tablet Hand lotion
Aspirin tablet Laundry detergent
Baking powder Lemon-lime soda
Baking soda Lemon juice
Bleach Milk
Club soda Mouthwash
Cola Shampoo
Contact lens solution Tea
Cream of tartar Toothpaste
Drain cleaning solution Vinegar
Fruit Fresh
Vitamin C tablet
Ginger ale Windex
6
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7
2009 Flinn Scientific, Inc. All Rights Reserved. 91564
Flinn ScientificTeaching Chemistry
).
a. What is the pH range in which the most basic form predominates?
b. Calculate the highest H
3
O
+
concentration (lowest pH) at which the indicator still exists in the In
form.
3. For one of the unknown acidbase solutions that were tested, explain why you chose the combination of indicators you did
to determine the pH value of the solution. What is the advantage of using multiple indicators, rather than a single indicator,
to determine the pH of a substance?
4. Construct a Results Table to summarize the properties of the unknowns.
a. Estimate the pH value of each unknown.
b. Classify each solution as acidic or basic.
c. Within each class of unknownsacids and basesarrange the solutions in order from least acidic to most acidic and
least basic to most basic, respectively.