Experiment 9
RED CABBAGE INDICATOR
Chemistry 51
INTRODUCTION:
Chemical indicators are weak organic acids or bases, dissolved in water, which change color if the+hydrogen ion concentration changes in a
solution. By definition, an Arrhenius acid is a substance that
base is a
- increases the concentration of hydrogen ion (H ) in a solution. An
+ Arrhenius
substance that increases the amount of hydroxide ion (OH ) in a solution. In water and in neutral solutions, the amounts of H and OH are the same.
The reaction below shows how hydrogen ion and hydroxide ion are formed from water.
HO
 H+ + OH-
In pure water and in neutral solutions, the amounts of hydrogen ion and hydroxide ion are very small.
+
An Arrhenius acid has hydrogen as the first element in its formula. When an acid dissolves in water, the acid releases H into the solution
and- the solution becomes acidic. An Arrhenius base has hydroxide ion in its formula. When a soluble Arrhenius base dissolves in water, it releases
OH into the solution and the solution becomes
basic.
Indicators are used to determine whether a solution is acidic, basic or neutral. Indicators impart
+
different colors depending on how much H or OH is in the solution.
The term pH is a+ simple way
of expressing how acidic (or basic) a solution is. The standard
pH scale runs from 0 to 14. If the pH is less than
+
7, the solution has more H than OH and the solution- is acidic.
The lower the pH, the more H is present and the more acidic
the solution is said to be.
+
If the pH is greater than 7, the solution has more OH than H and the solution is basic. The
higher
the pH, the more OH is present and the more basic
+
the solution is said to be. If the pH is equal to 7, the solution has the same amount of OH as H and the solution is said to be neutral.
You have already used indicators to predict whether a solution was basic or acidic. Litmus is an indicator that is blue in basic solution (pH >
7) but takes on a red color in acidic solution (pH < 7). Phenolphthalein is an indicator that is pink in solutions with pH values greater than 8, but is
colorless in solutions with pH values less than 8.
A buffer solution is a solution that will resist large pH changes upon the addition of an acid or base. Buffers will also maintain the pH when
they are diluted. Buffers resist pH change because there is an acid component to neutralize added base and a base componet to neutralize added acid.
There are many types of buffers, ranging from an acetic acid /acetate solution to natural buffers composed of amino acids and proteins. The most
common type of buffer in the chemistry laboratory is the combination of a weak acid like acetic acid and its conjugate base, an acetate salt. The
conjugate acid-base pair differ by an H+ ion.
In this experiment you will first investigate the change that happens when sodium hydroxide is added to water, then an acetic acid solution ,
and finally a buffer solution composed of acetic acid and its conjugate salt. It will be shown that a weak acid and, to a greater degree, the buffer
solution will resist a pH change when a base is added because the acid neutralizes the added base but dilution of the base with water will not cause
significant change. Next you will use an organic indicator extracted from red cabbage. You will examine the range of colors it imparts when put into
solutions with various pH values. In part B, you will record the colors of the indicators in solutions of known pH, or buffers. In doing so, you will obtain a
set of reference standards. In part C, you will investigate the acidity or basicity of various solutions by observing the color the indicator imparts in each
solution and comparing these colors to the reference standards you obtained in part B.
PROCEDURE :
Part A : Label three 50-mL beakers as "water", "acetic acid", and "buffer". Place 10 mL of distilled water in the
beaker labeled "water". Place 10 mL of 1.0 M acetic acid in the beaker labeled "acetic acid". Place 5 mL each of 1.0 M
acetic acid and 1.0 M sodium acetate in the beaker labeled "buffer" then mix well.
-Measure and record the pH of each solution using either pH paper or a pH meter. Your instructor will inform you
on which one to use.
-Add one drop of 1.0 M NaOH to each beaker. Measure and record the pH of each solution.
-Add a second drop of NaOH to each beaker. Measure and record the pH of each solution.
-Add an additional 18 drops (for a total of 20 drops; this is equilivant to 1 mL) to each beaker then measure and
record the pH the solution in each beaker.
-Next add 1 mL aliquots (portions) to each beaker, remembering to record the pH at each aliquot. Continue until a
total volume of 11 mL is added to each beaker.
Part B : Place about 10 - 20 drops of pH 1, 2, 4, 6, 7, 8, 10, 12, 14 standard buffers into separate wells in a spot
plate and label them with tape or a marker. To each well containing a buffer solution, add 0.5 mL (about 10 drops) of red
cabbage indicator. Thoroughly mix each solution then observe the color of the solution in each test tube. Record your
observations on the report sheet. Do not discard these solutions. Use them for reference for part B.
Part C : Pour about 10 - 20 drops of each solution listed on the report sheet in separate pre-labeled wells in a
spot plate. Next add 0.5 mL of cabbage indicator to each tube and mix well. Record your observations. Determine the
pH of each solution to the nearest pH unit and enter your value on your report sheet. Also indicate if each solution is an
acid, base, or neutral. Calculate the hydronium ion concentration in each solution.
Part D : In a spot plate using the first two rows, place in separate wells one piece each of zinc metal, copper
metal, a paper clip, and chalk (CaCO3). In the third well position below the two rows, place a small amount of sodium
bicarbonate also called baking soda (NaHCO3) in each well. To the first row you should add several drops of 6.0 M HCl
acid to each substance. Record you observations on the report sheet. Next add 6.0 M NaOH (a base) to each substance
in the second row. Record your observations.
Expt. 9
Name:
Report Sheet
Part A - A study of buffers
Amount of NaOH
pH in water
pH in acetic acid
1 drop NaOH
__________
______________
2 drops NaOH
__________
______________
20 drops NaOH
__________
______________
2 mL NaOH
__________
______________
3 mL NaOH
__________
______________
4 mL NaOH
__________
______________
5 mL NaOH
___________
______________
6 mL NaOH
___________
______________
7 mL NaOH
___________
______________
8 mL NaOH
___________
______________
9 mL NaOH
___________
______________
10 mL NaOH
___________
______________
11 mL NaOH
___________
______________
ph in acetic acid-acetate buffer
Part B - Standard Buffer Solutions with Red Cabbage Indicator
Buffer pH
Color of indicator
Acidic, basic, or neutral
1
_______________
____________________
2
_______________
____________________
4
_______________
____________________
6
_______________
____________________
7
_______________
____________________
8
_______________
____________________
10
_______________
____________________
12
_______________
____________________
14
_______________
____________________
+
[H3O ]
Part C - pH of Various Solutions
Solution
Color of indicator
DISH SOAP
______________
pH
_____
Acidic, basic, or neutral
____________________
BAKING SODA
______________
_____
____________________
TAP WATER
______________
_____
____________________
0.1M HCl
______________
_____
____________________
0.1M NaOH
______________
_____
____________________
VINEGAR
______________
_____
____________________
7-UP/SLICE
______________
_____
____________________
LEMON JUICE
______________
_____
____________________
Liquid antacid
______________
_____
____________________
BLEACH
______________
_____
____________________
AMMONIA
______________
_____
____________________
Part D: Reaction of Acids and Bases.
Table of observations:
Zinc metal
Copper
metal
Paper clip
(Al)
Chalk (CaCO3)
Sodium
bicarbonate
Addition
of HCl
Addition
of NaOH
GENERAL QUESTIONS:
1. Which is best at resisting pH change when a base is added, water, acetic acid, or buffer? ______
2. Which is worst at resisting pH change when a base is added? ________
3. Why does the buffer solution experience a large change in pH after 10 mL of base is added?
4. Why would it be useful to "buffer" an aspirin tablet? (Aspirin is acetylsalicylic acid).
5. Why is it a good idea to add the same amount of indicator to each of the test solutions in part "B"
as was added to your reference standards in part "B"?
6. What is the pH of a neutral solution? ____________
7.
What is the pH range of an acidic solution? __________________
8.
What is the pH range of a basic solution?
__________________
9. There are a large number of synthetic and natural indicators available, the pH where color
change occurs differs from indicator to indicator. Titration is a technique where a base is added to an
acid until the moles of acid equals the moles of base. Once all the moles of acid have been
consumed only moles of base remain in solution. Why would it be important to carefully choose your
indicator when doing an acid-base titration?
10. Do all metals react with acids? __________________
11. Zinc, Chalk, and baking soda only react with _________ but not with ____________.
12. Write an equation for the reaction of hydrochloric acid with zinc.
13. Write an equation for the reaction of hydrochloric acid with chalk.
14. Write an equation for the reaction of hydrochloric acid with baking soda.
15. Write an equation for the reaction of sodium hydroxide with zinc. Why is this equation different
from the others written above?