2606 Properties of weak and strong acids

1.0 INTRODUCTION

According to the Arrhenius definition, acids ionize in water to produce a hydronium ion (H3O+), and bases dissociate in water to produce hydroxide ion (OH-). An aqueous (water) solution that has a lot of hydronium ions but very few hydroxide ions is considered to be very acidic. If instead an aqueous solution has a lot of hydroxide ions but very few hydronium ions, it is considered to be very basic. Acids and bases are measured on a scale called pH. The pH is the negative of the log of the hydronium ion concentration.

pH = -log[H3O+]

pH ranges from less than 1 to 14. It lets us quickly tell if something is very acidic, a little acidic, neutral (neither acidic nor basic), a little basic or very basic. A pH of 1 is highly acidic, a pH of 14 is highly basic, and a pH of 7 is neutral.

pH indicators, litmus paper, and pH paper can be used to determine whether something is an acid or a base and the strength of its acidity or basicity. An indicator is a substance that turns a different color at a certain pH. Litmus paper is a form of an indicator. It is made by coating paper with the indicator litmus. Litmus is known to change color at a pH of about 7. Either red or blue litmus paper can be purchased. Blue litmus paper remains blue when dipped in a base, but it turns red when an acid touches it. Red litmus paper stays red when dipped in an acid, but turns blue when a base touches it.

Another way to more specifically determine an acid or base is through the use of pH paper. pH paper allows us to determine to what degree a substance is acidic or basic. When a substance is placed on pH paper, a color appears. The color is compared to a color chart showing the color the pH paper will turn at different pH values.

The use of a pH meter provides a more precise measurement of pH. The pH meter converts voltage readings from a pH electrode, an electrode constructed from a special composition glass which senses the hydrogen ion concentration. This thin, fragile glass is typically composed of alkali metal ions. The alkali metal ions of the glass and the hydrogen ions in solution undergo an ion exchange reaction, generating a potential difference.

1.1 Objectives

After completing this experiment, the student will be able to:

• Classify substances as weak acids or strong acids.
• Determine the pH values of several acid solutions using pH paper and a pH meter.
• Explain the effect on pH when adding a common ion to a weak acid solution.

2.0 SAFETY PROCEDURES AND WASTE DISPOSAL

3.0 CHEMICALS AND SolutionS

Chemical	Concentration	Approximate Amount	Notes
HCl solution	Between 1 M and 6 M	5 mL
HNO3 solution	Between 1 M and 6 M	5 mL
H2SO4 solution	Between 1 M and 6 M	5 mL
Acetic acid HC2H3O2 solution	Between 1 M and 6 M	5 mL
Acetic acid HC2H3O2 solution, standardized	1 M	25 mL	Record exact concentration
Oxalic acid dihydrate H2C2O4 . 2H2O	solid	1 g
Potassium hydrogen phthalate (KHP, HKC8H4O4)	solid	1 g
Sodium acetate NaC2H3O2	solid	2 g

4.0 GLASSWARE AND APPARATUS

pH meter (fragile!)	10 mL volumetric pipet and bulb
Several beakers, any size 150 mL or smaller	Two 100 mL volumetric flasks with stoppers
10 mL graduated cylinder	pH paper
50 mL graduated cylinder

5.0 PROCEDURE

5.1 Preparation and pH Measurement of Dilute Acid Solutions

1. Prepare 100 mL of the following solutions at approximately 0.1 M. You will need to use the dilution equation M1V1 = M2V2 to calculate how much volume to transfer from each stock solution provided. For the acids that are solid, use its molar mass to calculate how much mass to weigh. Use 150 mL or larger beakers.

2. Show your calculation for one of your solutions below.

Acid	Concentration of stock solution	Amount measured out for a 0.1 M solution
HCl solution
HNO3 solution
H2SO4 solution
Acetic acid HC2H3O2 solution, standardized
Oxalic acid dihydrate H2C2O4 . 2H2O	N/A
Potassium hydrogen phthalate (KHP, HKC8H4O4)	N/A

3. Use pH paper to estimate the pH of each solution and record in Table 1 found in 6.0 DATA RECORDING SHEET below.

5.2 Measuring How pH Changes with Concentration

1. In a clean, dry beaker, obtain approximately 25 mL of the 1.0 M Standardized Acetic Acid solution. Label the beaker “Stock Solution”.

2. Fill a 10 mL volumetric pipet with “Stock Solution” and drain into the sink. Using the rinsed volumetric pipet, transfer 10 mL of “Stock Solution” to a 100 mL volumetric flask, fill to the mark with laboratory water, and mix well. Label the flask “Dilution #1”.

3. Fill the 10 mL volumetric pipet with “Dilution #1” and drain into the sink. Using the rinsed volumetric pipet, transfer 10 mL of “Dilution #1” to a second 100 mL volumetric flask, fill to the mark with laboratory water, and mix well. Label this flask “Dilution #2”.

4. Record the concentration of “Stock Solution” in Table 2 found in 6.0 DATA RECORDING SHEET below, then measure and record the pH of each solution using a pH meter (pour some of each solution into a beaker to make the pH measurements).

5. The procedure below describes how to use a portable Flinn pH meter (model AP8673). If the pH meter is not measuring properly (unstable, or inaccurate), refer to the manual for troubleshooting. This same procedure may be applied to other brands of pH meter.

Warning: The pH meter probe (end or tip) is extremely fragile! Do not allow the tip to touch the flask or other solid object (except filter paper).

1. Check out a portable pH meter from the stockroom.
2. Remove the protective cap on the electrode. Clean any salt build-up off by rinsing with laboratory water.
3. Press the ON/OFF button once.
4. Rinse the electrode with laboratory water and blot dry with filter paper.
5. Immerse the electrode in the flask containing the analyte solution. Once the display stabilizes (approx. 1 min.), record the exact pH.
6. Remove the pH meter from the solution.
7. Repeat steps (d) – (g) after a certain volume of the titrant is addedif you are titrating.
8. When finished, rinse the pH meter electrode with laboratory water and blot dry with filter paper. Replace the cap and return the pH meter to the stockroom at the end of lab.

5.3 Measuring How pH Changes with Addition of a Common Ion

1. Use a 50 mL graduate cylinder to transfer 50 mL of “Dilution #1” to a clean dry beaker.

2. Measure its pH with the pH meter and record in Table 3 found in 6.0 DATA RECORDING SHEET below.

3. Dissolve 0.5 g solid sodium acetate in your 50 mL solution. In Table 3 below record the exact mass of sodium acetate that was added.

4. Measure its pH with the pH meter and record in Table 3 below.

5. Dissolve an additional 1.5 g solid sodium acetate to your 50 mL solution. In Table 3 below record the exact mass of sodium acetate that was added.

6. Measure its pH with the pH meter and record in Table 3 below.

6.0 DATA RECORDING SHEET

Last Name

First Name

Partner Name(s)

Date

Use the colors of the pH paper to estimate the pH of each solution from “5.1 Preparation and pH measurement of dilute acid solutions”. Then use your estimates to calculate the approximate H+ concentrations.

Table 1

Acid solution	Color of pH paper	pH	[H+]
HCl solution
HNO3 solution
H2SO4 solution
Acetic acid HC2H3O2 solution, standardized
Oxalic acid dihydrate H2C2O4 . 2H2O
Potassium hydrogen phthalate (KHP, HKC8H4O4)

1. Show your calculation of [H+] for one of your solutions below.

2. Given a concentration of 0.1 M, what is the predicted pH value for a monoprotic strong acid?

3. Analyze and comment on your data. For example, since all the acid solutions were prepared at the same concentration, do they all have the same pH value? If they do not, propose several reasons why not (do not include experimental error as a reason). Which acids had a pH close to the predicted value? Which acids did not? For each acid, propose reason(s) for their measured pH. Choose one acid and write a balanced chemical equation to show what happens when that acid dissolves in water.

Table 2

	Concentration of acetic acid	pH	[H+]	Actual [C2H3O2-]	Actual [HC2H3O2]
“Stock Solution”
“Dilution #1”
“Dilution #2”

1. Use the concentration and pH data to calculate the missing values in the data table above. Hint: Try using an ICE table.

2. In the space below, show the calculations for every value (except pH) in the row of the Table for “Dilution #1”.

3. Use your values from Table 2 to calculate the equilibrium constant for each solution and enter the results below.

	Calculated Equilibrium Constant
“Stock Solution”
“Dilution #1”
“Dilution #2”

4. Discuss how closely these constants agree with each other (Are they constant?)

Table 3

Concentration of acetic acid in “Dilution #1”	Additional grams of sodium acetate	pH	[H+]	Actual [C2H3O2-]	Actual [HC2H3O2]
	0

1. Calculate the missing values in Table 3. When sodium acetate is present, assume the “Actual [C2H3O2-]” is provided exclusively by sodium acetate. Remember to add together the two portions of sodium acetate (0.5 g + 1.5 g) when calculating the “Actual [C2H3O2-]” of the last solution!

2. Is the pH the same or different for all three solutions? Propose reasons for your observations.

3. Use your values from Table 3 to calculate the equilibrium constant for each solution and enter the results below.

	Calculated Equilibrium Constant
“Dilution #1”
“Dilution #1” With 0.5 g
“Dilution #2” With another 1.5 g

4. Discuss how closely these constants agree with each other (Are they constant?)

5. Do these equilibrium constants from Table 3 agree with the equilibrium constants calculated from Table 2? Should they agree?