Hydronium Ion concentration

Experiment 16
FV 2/11
HYDRONIUM ION CONCENTRATION
MATERIALS: 7 centrifuge tubes with 10 mL mark, six 25 x 150 mm large test tubes, 5 mL pipet, 50 mL beaker,
50 mL graduated cylinder, 1.0 M HCl, 1.0 M CH3COOH, CH3COONa, methyl violet and methyl
orange indicator solutions, test tube rack, calibrated pH meter (instructor use only).
PURPOSE:
The purpose of this experiment is to perform serial dilutions and use indicators to estimate the pH
of various solutions.
LEARNING OBJECTIVES:
1.
2.
3.
4.
5.
By the end of this experiment, the students should be able to demonstrate the
following proficiencies:
Prepare solutions by serial dilution.
Correlate the H3O+ ion concentration of a solution with its pH value.
Use indicators to estimate the pH of solutions of various acid concentrations.
Explain the common ion effect.
Calculate percent dissociation of a weak acid.
PRE-LAB: Read over the experiment and complete the pre-lab questions on p.E16-7 before the lab.
DISCUSSION:
In any aqueous solution, the product of the hydronium ion concentration and the hydroxide ion
concentration is equal to a constant, known as Kw. At a temperature of 25°C, this product is:
[
][ ]
K w = H 3O + ⋅ OH - = 1.0 x 10-14
(25° C)
If the [H3O+] concentration is altered, the [OH-] concentration changes so that the product of the two terms remains
1.0 x 10-14. In pure, distilled water at 25°C:
[H3O+ ] = [OH- ] = 10. x 10-7 M
(25° C)
To describe acidity in a simple way, the pH scale was adopted:
[
pH = - log H 3O +
]
Thus, for pure, distilled water at 25°C:
(
)
pH = - log 1.0 x 10-7 = 7.00
(25° C)
The practical range of pH in aqueous solutions at 25°C is from 0 to 14. For acidic solutions, the [H3O+] is greater
than that in distilled water and the pH is therefore < 7; for basic solutions the [H3O+] is less than that in distilled
water and the pH is > 7. Neutral solutions have [H3O+] = [ OH-] and therefore, the pH of such a solution is 7.00 at
25°C.
An acid-base indicator is a substance which changes color with changes of hydronium ion concentration. 1
This color change may occur near the neutral point, pH = 7, or at some other value of pH depending on the indicator.
A color change in an indicator does not necessarily indicate a transition from an acidic to a basic solution. For
example, an indicator such as thymol blue may change color as the pH goes from 1 to 3. Notice that both the initial
and the final solutions in this case are acidic.
1
For a discussion of pH and acid-base indicators, see “Chemistry”, Chang & Goldsby, pp. 741-743.
E16-1
PROCEDURE:
Part A. Preparation of Standard HCl Solutions of pH 0 to 5
Thoroughly rinse all glassware to be used in this experiment with distilled water.
1.
Rinse the 5 mL pipet with 1.0 M hydrochloric acid. Use the rinsed pipet to transfer 5.00 mL of the
hydrochloric acid solution to a 50 mL graduated cylinder (previously rinsed with distilled water). Add
sufficient distilled water to make the resulting solution 0.10 M HCl. Gently swirl the solution to mix.
2.
Rinse a large test tube with a small portion of the 0.10 M HCl and discard the rinses. Transfer the
remaining 0.10 M hydrochloric acid solution to the test tube and place it in the test tube rack in the position
marked "1" (indicating a pH of 1.00).
3.
Rinse the 50 mL graduated cylinder with distilled water. Rinse the 5 mL pipet with the 0.10 M HCl
solution which you just prepared. Use the rinsed pipet to transfer 5.00 mL of the 0.10 M HCl solution to
the 50 mL graduated cylinder. Add sufficient distilled water to make the resulting solution 0.010 M HCl.
After mixing, transfer this 0.010 M HCl solution to a large test tube and place it in the rack in the position
marked "2" (indicating a pH of 2.00). This process of generating a series of new solutions by successive
dilution of previously prepared solutions is called “serial dilution.”
4.
By serial dilution, prepare solutions of hydrochloric acid in which the concentrations are 1.0 x 10-3 M,
1.0 x 10-4 M, and 1.0 x 10-5 M (pH of 3.00, 4.00 and 5.00 respectively). As before, transfer each of these
to a clean large test tube and place in the rack.
5.
Finally, label one test tube "0" and place the original (undiluted) 1.0 M hydrochloric acid solution in it.
You should now have six large test tubes of hydrochloric acid solution varying in concentration from 1.0 M
to 1.0 x 10-5 M. These solutions have pH values very close to 0.00, 1.00, 2.00, 3.00, 4.00, and 5.00, since
hydrochloric acid is a strong acid that completely dissociates in water.
6.
Rinse a centrifuge tube with the 1.0 M HCl solution. After rinsing, fill the centrifuge tube to the 10 mL
mark with the 1.0 M HCl. Place this centrifuge tube in the position in the rack marked "0".
7.
Repeat this process for all the remaining solutions, first rinsing each tube with the solution it is to contain.
In the seventh tube, place an extra 10 mL sample of the pH 3.00 solution. Place each tube in that position
in the rack corresponding in number to the pH of the solution.
Part B. Color Changes of Various Indicators in Standard HCl Solutions
1.
Add one drop of methyl violet indicator solution to the four HCl solutions of pH 0.00, 1.00, 2.00, and 3.00
(which are in the centrifuge tubes).
2.
Starting with the extra tube containing the HCl solution of pH 3.00, add one drop of methyl orange
indicator solution to each of the three remaining solutions (pH 3.00, 4.00, and 5.00). Mix each solution
thoroughly, using a small glass stirring rod. Be sure to rinse the stirring rod between uses to prevent cross
contamination of the solutions.
3.
Record the observed colors of the solutions in the Data Section. (You may see different shades of the same
color; be as descriptive as possible.) The useful range for each indicator lies between the pH value for the
first and the last tube for that indicator. The color outside the useful range will be the same as the last color
in the useful range. Save these solutions for use as color standards for Parts C and D. Have your
instructor check these standards before proceeding.
E16-2
Part C. Color Changes of Various Indicators in Acetic Acid Solutions and Estimated pH Values
1.
Empty only the two centrifuge tubes containing the color standards for pH 0.00 and 1.00, and all the large
test tubes. Rinse all the tubes thoroughly with water. These tubes will be used for preparing and testing
various solutions of acetic acid.
2.
Serial dilution will again be used to prepare the acetic acid solutions. Starting with 5.00 mL of the 1.0 M
solution of acetic acid, prepare 50 mL of 0.10 M acetic acid. Use 5.00 mL of this solution to prepare 50 mL
of 0.010 M acetic acid. Then prepare 50 mL of a third solution of 0.0010 M acetic acid.
3.
Fill two of the empty centrifuge tubes to the 10 mL mark with the 0.0010 M acetic acid solution. Add
methyl violet to one centrifuge tube and methyl orange to the other.
4.
Compare these tubes with the color standards of Part B by looking down through the tubes at a piece of
white paper. (Indicator colors should always be compared directly with standards.) Record the colors in
the Data Section. Also, based on your experimental observations, estimate the pH of the solution to the
nearest half unit and record it in the appropriate space in the Data Section.
5.
In similar fashion, test 10 mL samples of the 0.010 M, 0.10 M, and 1.0 M acetic acid solutions with each of
the two indicators recording the colors and estimating the pH values. Calculate the corresponding
approximate concentrations of hydronium ion, H3O+, for both the HCl and acetic acid solutions based on
the pH values.
Part D. The Common Ion Effect
1.
Fill a centrifuge tube to the 10 mL mark with the 0.10 M acetic acid solution saved in part C and add one
drop of methyl orange indicator to it. Mix thoroughly. Record the color and estimate the pH based on the
color.
2.
In this solution dissolve as much solid sodium acetate, CH3COONa, as can be held on the end of a spatula.
Mix the solution until all of the solid dissolves. Note any change in color and estimate the pH of the
solution by comparison with the standards from Part B.
3.
Bring your solution to your instructor to be tested with a pH meter. How does your measured pH compare
to your estimate?
4.
After you clean up, start working on the calculations for this lab on page E16-5.
Clean up:
1.
Dispose of all aqueous solutions down the drain.
2.
Clean all glassware and return them to their original locations.
E16-3
Name _________________________________
Section _________________________
Partner _______________________________
Date ____________________________
DATA SECTION
Experiment 16
Part B. Observed Colors of Various Indicators in Standard HCl Solutions
pH
0.00
1.00
2.00
3.00
4.00
5.00
Methyl Violet
Methyl Orange
Instructor's Initials ____________
Part C. Observed Colors of Various Indicators in Acetic Acid Solutions
Acetic Acid
1.0 M
0.10 M
0.010 M
0.0010 M
Methyl Violet
Methyl Orange
Instructor's Initials ____________
Part C. Estimated pH Values of Hydrochloric Acid and Acetic Acid Solutions
Undissociated Acid
Concentration
Hydrochloric Acid
approximate pH
1.0 M
0.00
0.10 M
1.00
0.010 M
2.00
0.0010 M
3.00
Acetic Acid
calculated [H30+]
estimated pH
calculated [H30+]
Part D. The Common Ion Effect
Solution
observed color
estimated pH
0.10 M CH3COOH
0.10 M CH3COOH
containing CH3COONa
E16-4
measured pH
QUESTIONS
Experiment 16
1.
In this experiment the concentration of H3O+ ion in the solution of acetic acid could only be roughly
determined. More accurate values are given in the table below. Using these values, complete the table,
calculating any missing values. Clearly show the calculations for the first row in the table.
(Temperature = 25°C)
Initial conc. of
CH3COOH
Equilibrium
conc. of H3O+
1.00 M
0.00419 M
0.500 M
0.00296 M
0.200 M
0.00187 M
Equilibrium
conc. of
CH3COO-
Conc. of
undissociated
CH3COOH
Percent
dissociation
Ka
Calculations for 1.00 M initial conc. of CH3COOH:
2.
As a solution of acetic acid is diluted, do the calculations in the previous question show that:
a. the percent dissociation increases, decreases, or remains the same? Give a justification for your answer,
using your data and calculations. Explain why this occurs.
b. the dissociation constant, Ka, increases, decreases, or remains the same? Give a justification for your
answer, using your data and calculations.
E16-5
3.
Which of the two indicators cannot be used to estimate the pH of acetic acid of different concentrations?
Why not?
4.
a. In Part D, how does the addition of sodium acetate shift the equilibrium? Explain your answer. Show
any pertinent chemical reactions involved.
b. What effect will the addition of sodium acetate have on the pH of the solution? Should pH increase,
decrease, or remain the same? Explain.
c. Did the experimental results match your predicted change in pH? Explain any discrepancies.
E16-6
Name __________________________________
Section _____________________________
Date ___________________________________
PRE-LAB QUESTIONS
Experiment 16
1.
a. Write the balanced chemical reaction for the dissociation of HCl in water.
b. Write the balanced chemical reaction for the dissociation of acetic acid in water.
c. What is the difference between the two acids?
d. Which would have a lower pH, a solution of 0.10 M HCl or 0.10 M acetic acid? Explain your answer.
2.
Explain how a serial dilution is performed.
3. Using LeChatelier’s Principle, predict what should happen to pH when sodium acetate (CH3COONa) is added to a
solution containing acetic acid.
CH3COOH(aq) + H2O(l)
⇌
H3O+(aq) + CH3COO-(aq)
E16-7