Common Ion Effect

Introduction

In three demonstrations, students will observe a series of dramatic changes when common ions are added to solutions at equilibrium. The results illustrate the nature of equilibrium and Le Chatelier’s principle.

The set of three demonstrations includes:

1. Colorful Acid Eruptions

   Four acid solutions, two of which contain their conjugate bases, are added to cylinders. A white powder is added to each and the cylinders overflow with colorful foam.

2. Carbon Dioxide and pH

   A straw is inserted into a beaker filled with blue solution. Blowing into the straw produces bubbles that change the solution to yellow. Add a small amount of white powder and the solution reverts to blue. Blowing again into the straw changes the color of the solution back to yellow.

3. Weak Base Color Cycle

   A dropper full of yellow liquid is added to a beaker of water. Add a clear liquid and the solution turns orange. Add a white powder, stir, and the solution returns to yellow. Repeat the procedure using a dropper full of second clear liquid and the solution colors oscillate between yellow and orange.

Each demonstration includes an optional student worksheet that may be used for review or as an assessment tool by the instructor.

Concepts

• Weak acid
• Common ion effect
• Strong acid
• Le Chatelier’s principle
• Equilibrium

Background

Demonstration 1: Colorful Acid Eruptions
Four solutions, one a strong acid, one a weak acid, and two having the strong acid and the weak acid combined with their conjugate bases, are each added to a cylinder. Calcium carbonate and universal indicator are then added and each cylinder in turn fills with colored foam.

Demonstration 2: Carbon Dioxide and pH
Use a straw to blow bubbles into a blue solution until the solution color is pale yellow. Add a white powder and the solution returns to blue.

Demonstration 3: Weak Base Color Cycle
A dropper-full of yellow liquid is added to a beaker of water. Add a clear liquid and the solution turns orange. Add a white powder, stir, and the solution returns to yellow. Repeat the procedure using a dropper full of second clear liquid and the solution colors oscillate between yellow and orange.

Materials

Safety Precautions

Hydrochloric acid and acetic acid solutions are toxic and corrosive. Avoid contact with skin and eyes. “Rainbow-acid” universal indicator solution is an alcohol-based solution and is flammable. Avoid contact with flames or other ignition sources. Ammonium hydroxide solution is moderately irritable to skin and eyes. Ammonium chloride is slightly toxic by ingestion. Ammonium hydroxide solution is irritating to the skin and eyes and moderately toxic by ingestion or inhalation. Phenolphthalein solution is flammable and is moderately toxic. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Please consult current Safety Data Sheets for additional safety, handling and disposal information.

Disposal

Please consult your current Flinn Scientific Catalog/Reference Manual for general guidelines and specific procedures, and review all federal, state and local regulations that may apply, before proceeding. The waste solutions may be disposed down the drain with excess water according to Flinn Suggested Disposal Method #26b. Excess hydrochloric acid and acetic acid may be saved for future use or neutralized and disposed of according to Flinn Suggested Disposal Method #24b. The reaction products may be disposed of according to Flinn Suggested Disposal Method #26b.

Prelab Preparation

Demonstration 1: Colorful Acid Eruptions

1. Hydrochloric acid and acetic acid are supplied in the kit as 2 M solutions. Dilute 1:1 with distilled or deionized water to prepare 1 M solutions needed for the demonstration. For example, add 100 mL of 2 M hydrochloric acid to water and dilute to 200 mL with additional distilled water.
2. To save time in the presentation, pre-measure the amounts of solids needed for the demonstration. Weigh out 12 g of sodium chloride, 16 g of sodium acetate and 4 x 10-g samples of calcium carbonate in separate, labeled weighing dishes or small beakers.

Demonstration 2: Carbon Dioxide and pH

1. Add 600 mL of distilled water to an 800- or 1000-mL beaker.
2. Add 25 drops of bromthymol blue to the beaker. Stir.
3. If the solution is blue-green or yellow, add 0.1 M ammonium hydroxide drop by drop to the beaker, while stirring, until solution just turns bright blue.

Procedure

Demonstration 1: Colorful Acid Eruptions

1. Obtain 4 large hydrometer cylinders or tall-form beakers and place them on a large demonstration tray. Label the cylinders 1–4.
2. Using a graduated cylinder, add 100 mL of 1 M hydrochloric acid to cylinders 1 and 2.
3. Using a graduated cylinder, add 100 mL of 1 M acetic acid to cylinders 3 and 4.
4. Add about 20 drops (1 mL) of “rainbow acid” universal indicator to cylinders 1 and 3. Compare the color and pH of hydrochloric acid versus acetic acid.
5. Write equations for the ion-forming reactions of hydrochloric acid and acetic acid in water to give H₃O⁺ ions. Identify the “common ion” or conjugate base of each acid (chloride ion and acetate ion, respectively).
6. Add 12 g of sodium chloride, followed by about 20 drops of “rainbow acid” universal indicator, to cylinder 2. Mix thoroughly to dissolve.
7. Compare the color and pH of cylinder 2 with that in cylinder 1. What effect does adding chloride ion (its “common ion” or conjugate base) have on the pH of hydrochloric acid?
8. Add 16 g of sodium acetate, followed by about 20 drops of “rainbow acid” universal indicator, to cylinder 4. Mix thoroughly to dissolve.
9. Compare the color and pH of cylinder 4 with that in cylinder 3. What effect does adding acetate ion (its “common ion” or conjugate base) have on the pH of acetic acid?
10. Add 10 g of calcium carbonate to each cylinder 1–4.
11. Compare the amount of frothing and foaming and observe the rainbow indicator color changes in the four cylinders.
12. Relate the activity of the solutions to pH, the difference between strong and weak acids, and the “common ion” effect.

Demonstration 2: Carbon Dioxide and pH

1. Use the flexible straw to blow bubbles through the blue solution until the solution turns a yellow-green. Note: Use care so bubbles do not splash solution out of the beaker.
2. Use a spatula to add a tipful of sodium bicarbonate to the solution. Stir.
3. Repeat step 2 until the solution is blue-green to blue.
4. Again, use the straw to blow bubbles through the solution until the color changes back to yellow-green.
5. Repeat steps 1–3 until there are no further color changes.

Demonstration 3: Weak Base Color Cycle

1. Add 600 mL of distilled water to an 800- or 1000-mL beaker.
2. Add a dropper full of Alizarin Yellow R indicator solution to the beaker. Stir. The solution should be yellow in color.
3. Add 5 mL of 1.0 M ammonium hydroxide to the beaker and stir. The indicator color changes to orange.
4. Using a spatula, slowly add ammonium chloride to the beaker while stirring until the solution color returns to yellow.
5. Repeat steps 3 and 4.
6. The demonstration can be repeated, this time using phenolphthalein indicator solution. The solution will start out clear, then oscillate between red and light pink.

Student Worksheet PDF

13542_Student1.pdf

Teacher Tips

• This kit contains enough chemicals to perform the Colorful Acid Eruptions demonstration as written seven times: 750 mL each of 2 M hydrochloric acid and acetic acid, 280 g of calcium carbonate, 120 g of sodium acetate, 100 g of sodium chloride, and 50 mL of “rainbow acid” universal indicator.
• The use of a demonstration tray to catch any spillover (particularly from reactions 1 and 2) is strongly recommended. The reaction mixtures bubble and froth and a solid wall of foam may erupt out of the hydrometer cylinders.
• “Rainbow acid” universal indicator is an indicator solution that uses a combination of indicators to obtain a rainbow spectrum of colors for acid solutions having pH values between 1 and 7. Use the color chart that accompanies the indicator to estimate the pH of acidic solutions.
• This kit contains enough chemicals to perform the Carbon Dioxide and pH demonstration as written seven times: 25 g of sodium bicarbonate, 25 mL of bromthymol blue indicator solution, 10 mL of 0.1 M ammonium hydroxide and 14 Beral-type pipets.
• Try to add only enough sodium bicarbonate or carbon dioxide to just cause a color change. This will reduce the amounts of NaHCO₃ and carbon dioxide needed for each color change and will allow more color cycles to be observed.
• This kit contains enough chemicals to perform the Demonstration 3: Weak Base Color Cycle demonstration seven times: 25 g of ammonium chloride, 50 mL of 1 M ammonium hydroxide, 20 mL of 1% phenolphthalein solution, 20 mL of 0.1% Alizarin Yellow R solution and 14 Beral-type pipets.
• Alizarin Yellow R indicator is an acid–base indicator that is yellow when the pH < 10.2, red when the pH > 12.2 and various shades of orange in between. Other indicators that change color between the range of pH 8–10 or pH 10–12.5 can also be used, such as thymolphthalein (colorless to blue, pH 9.9–10.5) and alizarin (red to purple, pH 11.0–12.4).

Sample Data

Demonstration 1: Colorful Acid Eruptions

Effect of Common Ion Based on these observations, are both reactions reversible? If not, why?

Reversible reactions show the common ion effect. The additions of chloride ion has no effect on the hydrochloric acid ionization reaction. This means that the hydrochloric acid ionization reactions is irreversible. The addition of sodium acetate to the acetic acid solution reduces the hydronium ion concentration so the reaction shifts to the left side of the acetic acid ionization reaction.

Demonstration 2: Carbon Dioxide and pH
Carbon dioxide dissolved in water is a weak acid, establishing equilibrium with its products, the hydronium ion and the bicar-bonate ion. The solution is initially acidic. Bromthymol blue is yellow in an acidic solution, blue-green in a neutral solution, and blue in a basic solution. What is the color change when bubbles are blown into the solution? Why does it occur?

Blue to yellow. Blowing into the solution dissolves more CO2. This forces the equilibrium to shift to the right, reestablishing equilibrium at an increased concentration of hydronium ions.

Explain the color change when powder, sodium bicarbonate is added to the solution.

The increase in bicarbonate ion causes a shifts to the left, consuming the hydronium ions until equilibrium is reestablished at higher pH.

Demonstration 3: Weak Base Color Cycle
“Ammonium hydroxide” is a concentrated solution of ammonia in water. Ammonia (NH₃) is a weak base, reacting with water to form ammonium ions and hydroxide ions (Equation 1). Explain why the solution changes color when the solid, ammonium chloride, NH₄Cl, is added to the beaker.

Increased ammonium ion leads to decrease in hydroxide ion as the reaction shifts to the left to reestablish equilibrium. Lower pH causes indicator to change color.

Discussion

Demonstration 1: Colorful Acid Eruptions

The difference between strong and weak acids is confusing to many students. Some students assume that the difference arises solely due to pH. Others attribute the difference to concentration. This demonstration compares the pH, activity, and “common ion” effect in both hydrochloric and acetic acid solutions to illustrate the properties of strong versus weak acids.

Hydrochloric acid and acetic acid are Brønsted acids—they ionize in water to produce hydrogen ions (H₃O⁺) and their conjugate bases, chloride ion and acetate ion, respectively (Equations 1 and 2).

Comparing the pH of these two acids indicates that the amount of hydrogen ions produced in the two solutions is very different—there are more H₃O⁺ ions present in the hydrochloric acid solution than in acetic acid. In the hydrochloric acid solution, all of the HCl molecules undergo ionization to form H3O+ ions. In acetic acid, however, only a few H₃O⁺ ions are produced (most of the CH₃CO₂H molecules are not ionized). Note that this comparison—the relationship between pH and the degree of ionization of hydrochloric acid versus acetic acid—is only valid because their initial concentrations are the same (1 M).

The activity of the two acids with calcium carbonate, a strong base, reinforces the pH comparison. The rate of reaction of calcium carbonate with an acid (Equation 3) depends on the concentration of hydrogen ions in solution. The foam produced in this reaction is due to carbon dioxide gas mixing with water and calcium carbonate powder. The amount of foaming and the rate at which the foam rises are dramatically different for 1 M hydrochloric acid versus 1 M acetic acid. The reaction with hydrochloric acid is significantly faster, suggesting again that the concentration of H₃O⁺ ions is greater than in acetic acid. A “rainbow spectrum” of indicator color changes is observed as the pH changes and the reaction proceeds. The effect of chloride ion and acetate ion on the pH and reactivity of hydrochloric acid and acetic acid, respectively, further distinguishes the behavior of strong versus weak acids. Adding chloride ion to hydrochloric acid does not change either the pH or the activity of the acid solution. This suggests that the reaction shown in Equation 1 takes place in one direction only—ionization of hydrochloric acid is irreversible. Adding acetate ion to acetic acid, however, increases the pH of the solution from 2 to almost 5 and drastically slows down its reaction with calcium carbonate. Both of these observations suggest that the hydrogen ion concentration in the mixed acetic acid/sodium acetate solution is 100–1000X lower than in acetic acid itself. The reaction shown in Equation 2 is thus effectively reversed in the presence of acetate, the “common ion.” Ionization of acetic acid is reversible (Equation 4) and the equilibrium constant for this reaction is very small (approx. 10^–5). Demonstration 2: Carbon Dioxide and pH

The initial color of distilled water with bromthymol blue indicator is ususally blue-green or pale yellow due to the presence of dissolved carbon dioxide from the air. Carbon dioxide dissolved in water is a weak acid, establishing equilibrium with its products, the hydronium ion and the bicarbonate ion. The solution is acidic. Bromthymol blue is yellow in an acidic solution (pH < 6.0), blue-green in a neutral solution (pH = 6.0–7.6) and blue in a basic solution (pH > 7.6). When ammonium hydroxide is added to the beaker, the hydronium ions are neutralized or consumed. Removing the H₃O⁺ ions causes the equilibrium to shift to the right, reducing the amount of dissolved carbon dioxide. This continues until all the carbon dioxide is consumed. The solution becomes basic and the solution color is blue.

When carbon dioxide is blown through the straw into the solution, the amount of carbon dioxide dissolved in solution increases. The reaction shifts to the right, producing hydronium ions. As these are consumed by the hydroxide ions, the reaction continues consuming carbon dioxide until all the hydroxide ions are neutralized. The equilibrium then is reestablished and the solution turns the acidic yellow color.

Adding the bicarbonate ion causes the reaction to shift to the left until the hydronium ions are consumed. This results in a color change to “various shades of green in the transition range (pH 6.0–7.6).

Demonstration 3: Weak Base Color Cycle
“Ammonium hydroxide” is a concentrated solution of ammonia in water. Ammonia (NH₃) is a weak base, reacting with water to form ammonium ions and hydroxide ions (Equation 6). When ammonium hydroxide is added to the beaker, the solution becomes very basic and its color changes from yellow (pH < 10.2) to orange (pH > 12.2). When ammonium chloride is added, the increase in the ammonium ion concentration “upsets” the equilibrium, causing a shift to the left to reestablish the balance. As hydroxide ions are consumed, the pH dips below 10.2 and the indicator color reverts to its original yellow color. Add more ammonium hydroxide and the reaction shifts right and the color goes to orange.

References

Shakhashiri, B. Z. Chemical Demonstrations: A Handbook for Teachers in Chemistry; University of Wisconsin: Madison, 1989; Vol. 3, p 155.