CHEM 1030L
Fundamentals of Chemistry Laboratory
Acids, Bases, Buffers, and pH
Background
Acids & Bases
Acids are compounds that donate protons (H+) in solution and bases accept protons in solution. Some
acids that we have used in lab this semester are hydrochloric acid (HCl) and nitric acid (HNO3). Acetic
acid (CH3COOH) is a common weak acid that can probably be found in your home. It is the acid that
gives vinegar its sour taste. Some common household bases are ammonia, detergents, and ovencleaning products. Sodium hydroxide (NaOH) is an example of a common base that we have used in
lab.
The pH scale
The pH scale is a tool used to describe the relative acidity, basicity, or neutrality of a solution. The pH
scale goes from 0 to 14, where values below 7 are considered acidic, values above 7 are basic and a pH
value equal to 7 is neutral.
Acidic
0
1
2
3
Neutral
4
5
6
7
Basic
8
9
10 11 12 13 14
Mathematically, the pH of a solution is equal to the negative log of the H+ ion concentration:
pH = -log[H+]
So, if [H+] = 1.00 x 10-4, then the pH = -log(1.00 x 10-4) = 4.00 . . . and the solution is acidic.
If [H+] = 2.50 x 10-11, then the pH = -log(2.50 x 10-11) = 10.6 . . . and the solution is basic.
The pH of a solution can be calculated using the equation above if the proton concentration is known
or if it can be measured. Common tools used to measure pH values are pH meters, pH paper, and
universal indicator. pH paper can be used to estimate the pH of a solution, but the pH meter provides a
very accurate pH measurement. The instructor will demonstrate the proper use and handling of the
pH meters. Another, more colorful way, to estimate the pH of a substance is by the use of an indicator.
Indicators are chemical compounds that change colors when the pH changes. Some common indicators
are listed below along with the respective pH range in which they change color.
Indicator
methyl orange
phenol red
phenolphthalein
alizarin yellow
pH range of Color
Change
3.1 - 4.4
6.4 - 8.0
8.1 - 9.6
10.1 - 12.0
Color in Acid Range
(low pH)
red
yellow
colorless
yellow
Color in Base Range
(high pH)
yellow
red
pink
orange-red
Since most indicators have a specific pH range in which they change color, a universal indicator is
often used to determine the pH of substance with unknown acidity or basicity. The universal indicator
is a solution that contains several indicators which allows for a wide range of pH determinations.
Revised 10/29/2015 – T. Ekman
Buffers
A buffer is a solution that can maintain a relatively constant pH when limited amount of acid or base
are added to it or when the solution is diluted. In other words, a buffer is a chemical system that resists
change in pH. Buffered solutions play a very important role in controlling the pH of all body fluids. For
example, a buffer system maintains the pH of blood between 7.35 and 7.45. If the pH of the blood goes
above or below that range, it can destroy blood cells.
Buffers maintain the pH of a solution by reacting with small amounts of acids or bases. A typical buffer
may be composed of a weak acid and its salt. The weak acid reacts with added base and the anion of
the salt reacts with added acid. This way, compounds that would otherwise increase acidity or basicity
of bodily fluids are largely neutralized by the buffer system and a specific pH is maintained. For
example, in blood the bicarbonate buffer system includes a weak acid, carbonic acid (H2CO3), and the
anion of carbonic acid, bicarbonate (HCO3-). When strong base (OH-) is added, it reacts with the weak
acid to produce the bicarbonate ion and water
H2CO3 + OH- —> HCO3- + H2O
When strong acid enters the blood, the H+ reacts with the HCO3- anion and forms carbonic acid
HCO3- + H+ —> H2CO3
In this way, any strong acid (H+) is removed and a much weaker acid (H2CO3) is produced. If strong
base is added, it is removed and a much weaker base (HCO3-) is produced. In either case, the pH will
shift a small amount because of the increased amount of weak acid or base. However, the change is
much less than would have occurred if the strong acid or base had remained.
Procedures
Part A – Measurement of pH of Household Products
If the instructor allows students to work in groups of two, it is important that all steps be carried out
simultaneously and that each student observes the results obtained by the lab partner.
1. Using universal indicator
a. Use a clean, dry 50-mL beaker to obtain approximately 10 mL of one of the household products
from the back bench. (If working in pairs, each student should obtain a different household
product.)
b. Add 2-3 drops of universal indicator to the sample and stir with a glass rod.
c. Compare the color of the sample to the reference standards. The reference standard that most
closely matches the color of your sample indicates the pH of your sample. (Hint – Hold a sheet of
plain white paper behind or below the beaker to more easily determine the color.) Estimate the
pH and record.
2. Using pH paper
a. Place a strip of pH paper on a clean, dry watch glass.
b. Stir the solution from Step 1 with a clean glass stirring rod. Remove the stirring rod and touch
the tip of it to the pH paper so that a drop of the sample wets the paper. The sample will react
with indicator in the paper and cause a color change.
Revised 10/29/2015 – T. Ekman
c. Compare the color of the moistened pH paper to the reference chart on the pH paper vial. Match
the color and record the estimated pH.
3. Using the pH meter
a. The instructor will demonstrate the proper use of the pH meter.
b. Once the pH meter is calibrated, place the pH probe in the above sample and record the pH.
4. Repeat steps 1-3 until all six rows of the data sheet are completed.
Part B – Buffers
If working in pairs, each student will work with two of the solutions mentioned below and compare
results with the lab partner.
1. Effect of added acid
a. Place 10 mL of each of the following solutions into separate clean, dry 50-mL beakers.
 distilled water
 0.1 M NaCl
 carbonate buffer
 phosphate buffer
b. Use a pH meter to determine the initial pH of each sample. Record.
c. Add 5 drops of 0.1 M HCl to each sample. Stir and record the pH of each solution again.
d. Add another 5 drops of 0.1 M HCl to each solution. Stir and record the pH of each solution.
e. Clean the beakers by pouring the contents down the drain. Rinse with three small portions of
tap water followed by three small portions of distilled water. Dry with paper towel.
2. Effect of added base
a. Prepare four new solutions as described in Step B.1a, above.
b. Use a pH meter to determine the initial pH of each sample. Record.
c. Add 5 drops of 0.1 M NaOH to each sample. Stir and record the pH of each solution again.
d. Add another 5 drops of 0.1 M NaOH to each solution. Stir and record the pH of each solution.
e. Clean the beakers by pouring the contents down the drain. Rinse with three small portions of
tap water followed by three small portions of distilled water.
Revised 10/29/2015 – T. Ekman