CHEM 2124
General Chemistry II Laboratory
Experiment 1 - Indicators and pH of Aqueous Solutions and Household Products
INTRODUCTION
An acid is a substance which dissociates to produce hydrogen ions, H+, when dissolved in aqueous
solution. Once in solution, the H+ion, which is simply a proton, immediately combines with water to form
the hydronium ion, H3O+. So when aqueous H+ appears in a chemical equation, it is understood that
the actual species exists as H3O+. An acid is classified as strong or weak depending on the extent to
which the molecules of the acid dissociate into H+ and its anion, A-. A strong acid dissociates completely
in water (e.g., HCl dissociates virtually 100% into H+ and its anion, Cl-), while a weak acid dissociates only
partially and forms very little H+ (e.g., only about 3% of the dissolved molecules of acetic acid, CH3COOH,
dissociate into H+ and the acetate anion, CH3COO-). Weak acid dissociation can be represented as a
generic reversible reaction:
HA(aq) ⇔ H+ + Awhere HA is the weak acid and A- is the anion, or conjugate base, of HA. For this reversible reaction, an
equilibrium constant can be defined:
The equilibrium constant is called the acid dissociation constant or acid ionization constant. If the
undissociated acid (HA) is favored and the acid is weak, Kais measurable and will be much less than one.
For strong acids, the dissociated products, (H+ and A-) are so strongly favored, that [HA] approaches zero
and Ka approaches infinity. Thus strong acids, of which there are very few, do not have values of Ka
associated with them.
For convenience a log scale is often used to express very small quantities. For most aqueous solutions of
interest, the concentration of H+is small, ranging from approximately 10-1 to 10-12 M. Since it is also more
convenient to deal with positive numbers, the negative value of the log (–log, symbolized by “p”) of [H+] is
used to express the concentration of H+. Thus the definition of pH is:
pH = –log[H+] , and [H+] = 10-pH
(2)
As noted earlier, it is always understood that H+ is in fact present in solution as the hydronium ion, H3O+.
Notice that larger concentrations of H+ mean smaller values of pH.
Other –log scales can be defined in a similar fashion. One such scale is pOH.
pOH = –log[OH-]
In aqueous solutions, the concentration of H+ and OH- are always related by the dissociation constant of
water, Kw, which arises from the reaction called the autoionization of water.
H2O ⇔ H+ + OHIn this equilibrium, the molecular H2O is heavily favored. In fact the concentrations of the ions in pure
water are [H+] = [OH-] = 1 x 10-7M at 25°C. Thus the value of Kw is given by
–
-14
Kw = [H+][OH ] = 1 × 10
o
at 25 C.
(3)
In an aqueous solution that isn’t acidic or basic, the concentration [H+] is equal to that of [OH-], so the
solution is neutral, the same as pure water. Since this occurs at a concentration of 1 × 10-7M for both ions.
The values of pH and pOH are both 7 in a neutral solution.
pH = –log[H+] = –log(1 × 10-7) = 7
pOH = –log [OH-] = –log(1 × 10-7) = 7
The following relationships regarding aqueous solutions should therefore be understood.
pH
[H+]
Solution
less than < 7
> 1.0 × 10-7
Acidic
greater than > 7 < 1.0 × 10-7
Basic
equal to = 7
= 1.0 × 10-7
Neutral
Note also that if one takes –log of both sides of the equation for Kw (equation 3), the result is:
pKw = pH + pOH = 14
So in any aqueous solution, for which it is always true that Kw= 1 × 10-14 at 25°C, the sum of pH and pOH
of the solution must equal 14. Thus an acidic solution with pH = 4 must have a pOH of 10.
Part A - pH of Weak vs. Strong Acid
When a solution of acid is prepared, its initial molarity (before any dissociation of the acid molecules
occurs) is given by the moles of acid added divided by the liters of solution. For a strong acid, the
concentration of [H+] will be the same as the initial molarity of the acid because all the acid dissociates
into H+ and the anion A-. Because [H+] = [HA] in this case, thus the pH of a strong acid can be predicted
from the molarity of the acid. You will measure the pH of a strong acid to verify this relationship.
For a weak acid with much less than 100% dissociation, [H+] is much less than the initial [HA]. You will
measure the pH of a weak acid of known initial [HA]’initial’ Then [H+] can be calculated using equation (2).
Since [A-] is the same as the [H+] just calculated, and [HA] at equilibrium can be calculated from [HA]initial−
[H+], the value of Ka for the acid can be calculated from these values using equation (1).
Part B – Preparation of Buffer Solutions and Indicator Determinations
Many naturally occurring dyes in solution change color with the pH of the solution. You will add a
number of these dyes to buffer solutions of known pH, and observe how the dye color changes with pH.
In this part, you will need to prepare various dilutions which will serve as buffer solutions of pH 2
through 12.
Part C – pH of Common Household Products
You will test the pH of a number of household products using the various indicators.
Part D – Verification of pH Using pH Meter
In this part you will use a digital pH meter to accurately determine the pH of the various household
products for which the pH was estimated in part C.
PROCEDURE
Part A – pH of Weak vs. Strong Acid
1. Collecting data with the pH electrode: To obtain the pH of a solution, place the pH electrode in the
solution and wait for 15-30 seconds for the reading on the main screen to stabilize. Before placing the
pH electrode in a new solution it must be rinsed thoroughly with DI water and blotted. Strong bases
are hard to rinse off, take extra care in cleaning the electrode after measuring a basic solution (pH >
7). It may be necessary to dab the pH electrode with a Kimwipe to clean a basic solution off. Note: pH
meters can be unpredictable. If you feel you are not getting reasonable readings, check with
your instructor.
2. Obtain about 400 mL of de-ionized water in a 600 mL beaker for use in the dilutions described below.
Label three 150 mL beakers: (1) 0.1 M stock; (2) 1:10 dilution; and (3) 1:100 dilution.
3. Obtain approximately 15 mL of the 0.1 M stock solution of HCl in the (1) labeled beaker, take it to your
bench, and prepare a 1:10 dilution of this solution by placing 5.0 mL of it into a 100 mL graduated
cylinder and diluting it with DI water to the 50.0 mL mark. Transfer the diluted solution to the (2)
labeled 150 mL beaker.
4. Prepare a 1:100 dilution of the acid by taking the solution prepared in step 3 (the 1:10 dilution) and
further diluting it by placing 5.0 mL of it into a 100 mL graduated cylinder and diluting it to the 50.0
mL mark with DI water. Transfer this new solution to the (3) labeled 150 mL beaker.
5. Measure and record the pH of all three solutions of HCl in the results section. Make sure to also record
the exact molarity of the stock solution given on the bottle.
6. Pour the HCl solutions into the waste bottle. Rinse the three labeled beakers with tap water and then
DI water and use them to repeat steps 3-5 with acetic acid, 0.1 M CH3COOH, stock solution filling in
the table in the results section.
Part B – Preparation of Buffer Solutions and Indicator Determinations
1. Preparation of Buffer Solutions in the Acid Range, pH 2 to pH 6.
a. Place approximately 100 mL of distilled water which has been boiled into a large clean beaker.
b. Fill one test tube about 2/3 full with stock 0.01M HCl solution. The [H+]=0.01 M for 0.01M HCl and
the pH of this solution is 2. This should have been verified in Part A.
c. Prepare the pH 3 solution by making a tenfold dilution of the pH 2 solution. Use the 10-mL
graduated cylinder to measure 9 mL of the boiled distilled water and add 1 mL of the pH 2
solution from the stock bottle to the distilled water.
d. Fill one test tube about 2/3 full with the pH 3 solution.
e. Prepare the pH 4 solution by taking 1 mL of the pH 3 solution and adding it to 9 mL of boiled
distilled water. Use the cleaned and rinsed 10-mL graduated cylinder for this.
f. Fill one test tube about 2/3 full with the pH 4 solution.
g. Repeat the process using 1 mL of the pH 4 solution and 9 mL of boiled distilled water to make a
test tube of pH 5 solution.
h. Repeat the process using 1 mL of the pH 5 solution and 9 mL of boiled distilled water to make a
test tube of pH 6 solution.
2. Preparation of Neutral 7 pH Buffer Solution
a. The boiled distilled water is pH 7. Fill a test tube 2/3 full with this.
3. Preparation of Buffer Solutions in the Base Range pH 8 to pH 12.
a. Fill one test tube about 2/3 full with stock 0.01M NaOH solution. The [OH-]=0.01M for 0.01M
NaOH and thus the [H+] = 1x12-12M. Therefore, the pH of this solution is 12.
b. Prepare the pH 11 solution by making a tenfold dilution of the pH 12 solution as you did in step
1.c. above. (1 mL of pH 12 solution and 9.0 mL of boiled distilled water)
c. Fill one test tube about 2/3 full with the pH 11 solution.
d. Continue this process to make 1 test tube each of pH 10 solution, pH 9 solution, and pH 8 solution.
4. Indicator Determinations
a. To each tube in your set of test tubes (which contain solutions of pH 2 through 12, add 2-4 drops
of your assigned indicator solution.
b. Gently swirl all of the test tubes and note the color of each tube in the proper place of the results
section.
c. Check with the instructor to verify the results of you test tubes.
d. Place you set of test tube with indicator solution on the front bench and mark what indicator is
present.
e. Using the various sets of test tubes on the front bench, fill in Table B in the data sheet section.
Part C – pH of Common Household Products
1. Choose any four of the household products from the front bench. For liquid samples, use about 3mL
in each test tube and for solid samples, place a scoopula tip full into each test tube and add 3mL
distilled water.
2. Add 3mL of the first household product to five separate test tubes.
3. Add 3mL of the second household product to five more test tubes.
4. Continue this way until you have 5 test tubes of all 4 different household products. (20 total test tubes)
5. For household product #1, add 2-4 drops of each of the 5 indicators to the five test tubes. You should
have a red cabbage tube, a phenolphthalein tube, etc. for product #1.
6. Continue in this manner for all five indicators on all 4 household products.
7. Gently swirl all of the test tubes and note the color of each tube in the proper place of the results
section.
8. Using the data from part B and part C, estimate the pH of each of the six household products.
Part D – Verification of pH Using pH Meter
1. Place about 15-20 mL of household product #1 (the same one you used in part C) into the 50mL
Erlenmeyer flask. (If you are using a solid household product, place a “large” scoopula tip full of solid
into 20 mL of distilled water for your sample).
2. Measure and record the pH (using the pH meter) of household product #1 in the results section.
3. Pour out the sample and rinse the 50mL Erlenmeyer flask with distilled water.
4. Repeat steps 1-3 for the three remaining household products you tested in part C.
CHEM 2124
General Chemistry II Laboratory
Lab #1 – Indicators and pH of Aqueous Solutions and Household Products
Name
Date
Pre-Lab Questions – MUST be completed prior to lab.
PL1. Explain the difference between a strong acid and a weak acid in terms of ionization.
PL2. What is the pH of an acid that has a [H+] of 1.5x10-6M. Please report your answer to the correct
number of significant figures!
PL3. A clear liquid is found on the benchtop of a laboratory and is to be disposed of properly. Describe
two different methods by which you could determine the pH of this clear liquid.
PL4. Explain what is meant by the term “autoionization of water”.
Part A – pH of Weak vs. Strong Acid
1. Show a sample calculation of the H+ molarity from the pH values and complete the table below.
HCl
(1) Stock Solution
(2) 1:10 Dilution
(3) 1:100 Dilution
(1) Stock Solution
(2) 1:10 Dilution
(3) 1:100 Dilution
Molarity of acid:
pH:
H+ Molarity, [H+]
Acetic Acid
Molarity of acid:
pH:
H+ Molarity, [H+]
2. Calculate Ka for the CH3COOH using the [H+] and [CH3COOH]initial for each of the three solutions in the
data table.
Stock Solution:
1:10 Dilution:
1:100 Dilution:
3. Discuss how the data show that HCl is a strong acid and CH3COOH is a weak acid.
4. Explain how it is possible that a solution of a weak acid can have a lower pH than a solution of a strong
acid.
Part B – Preparation of Buffer Solutions and Indicator Determinations
Fill in the table below with the colors observed for each buffer solution and each indicator.
Cabbage Juice
Phenolphthalein Methyl Red
Bromcresol
Green
Methyl Orange
pH 2
pH 3
pH 4
pH 5
pH 6
pH 7
pH 8
pH 9
pH 10
pH 11
pH 12
Part C – pH of Common Household Products
1. Fill in the table below with the colors observed for each household product and each indicator as well
as the approximate pH of the household product.
Name of Product Cabbage
Used
Juice
Phenolphthalein
Methyl
Red
2. Which of the products are acidic and which are basic?
Bromcresol
Green
Methyl
Orange
Approximate
pH
Part D – Verification of pH Using pH Meter
1. Fill in the table below with the pH values obtained via pH meter for each of the six household products.
Name of Product Used
Approximate pH from Part C
Actual pH from pH Meter
2. Calculate the %EDA for pH for each of the six household products. Assume that the actual pH is that
obtained from the pH meter. Since pH is based on a logarithmic scale, you cannot directly calculate
the %EDA using pH values. Rather you will first have to convert all of the pH values to [H+] values and
them plug the [H+] values into the %EDA equation. As a reminder,
%EDA = observed – actual x 100
actual
Name of Product Used
Observed [H+] from
Part C
Actual [H+] from
Part D
%EDA
3. If any of your %EDA values are greater than 10%, explain why you think the %EDA is so high for that
particular household product.