Materials Included In Kit

Additional Materials Required

Safety Precautions

Ammonium chloride, potassium chloride, sodium carbonate, sodium phosphate and zinc sulfate are slightly toxic by ingestion and are body tissue irritants. Universal indicator solution is an alcohol-based flammable liquid. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Remind students to wash their hands thoroughly with soap and water before leaving the laboratory. Please review 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. Solutions may be rinsed down the sink. Leftover solid salts may be placed in the trash, according to Flinn Suggested Disposal Method #26a.

Lab Hints

• Enough materials are provided in this kit for 5 classes of 30 students working in pairs. This laboratory activity can reasonably be completed in one 50-minute class period. The prelaboratory assignment may be completed before coming to lab.
• If desired, the instructor can prepare the water for the entire class by heating and cooling the distilled water before class begins.
• Remind students to use clean spatulas and clean glassware for each salt. Any residue or contamination can alter their results.
• Additional salts can be tested (e.g., sodium chloride, sodium acetate, aluminum chloride hexahydrate, sodium phosphate monobasic, sodium bicarbonate).
• The salts described in the background can be demonstrated as an engaging presentation by the instructor before the lab occurs.

Teacher Tips

• A video of this activity, Hydrolysis of Salts, is available at www.flinnsci.com as part of the Flinn Scientific Best Practices for Teaching Chemistry teacher resources videos.
• Another fun activity is the Hydrolysis of Salts—Acidic, Basic or Neutral?–Chemical Demonstration Kit, Flinn Catalog No. AP6187.

Further Extensions

Additional Salts that can be Tested (from Lab Hints)

Acidic
Al(H₂O)₆ ³⁺(aq) + H₂O(l) → [Al(H₂O)₅OH]²⁺(aq) + H₃O⁺(aq)
H₂PO₄^–(aq) + H₂O(l) → HPO₄ ^2–(aq) + H₃O⁺(aq)
Basic 
C₂H₃O₂^–(aq) + H₂O(l) → HC₂H₃O₂(aq) + OH^–(aq)
HCO₃^–(aq) + H₂O(l) → H₂CO₃(aq) + OH^–(aq)

Answers to Prelab Questions

1. Define hydrolysis.

   Hydrolysis of a salt is the reaction of the salt with water or its ions.

2. Define Lewis acid and Lewis base.

   A Lewis acid is a substance that accepts an electron pair. A Lewis base is a substance that donates an electron pair.

3. Sodium bicarbonate is added to water. Three drops of universal indicator are added to the solution and the solution turns purple, indicating the solution is basic. Write the balanced net ionic equation below to show the formation of OH–.

   HCO₃^–(aq) + H₂O(l) → H₂CO₃(aq) + OH^–(aq)

Sample Data

Answers to Questions

1. For each salt that was acidic, write the balanced net ionic equation to show the formation of H₃O⁺.

   NH₄⁺(aq) + H₂O(l) → H₃O⁺(aq) + NH₃(g)
   Zn²⁺(aq) + H₂O(l) → H₃O⁺(aq)+ [Zn(H₂O)₄(OH)]⁺(aq)

2. For each salt that was basic, write the balanced net ionic equation to show the formation of OH^–.

   CO₃^2–(aq) + H₂O(l) → HCO₃^–(aq) + OH^–(aq)
   PO₄ ^3–(aq) + H₂O(l) → HPO₄ ^2–(aq) + OH^–(aq)

Introduction

View beautiful color changes when salts are added to water in this fun lab activity! Watch as the hydrolysis of salts changes the solutions to vibrant colors!

Concepts

• Salt hydrolysis
• Acids and bases
• pH
• Lewis acids and bases

Background

Acidic and basic properties of aqueous solutions depend on the concentrations of hydrogen ions [H⁺] and hydroxide ions [OH^–]. Water (the solvent in an aqueous solution) dissociates to a small extent into hydrogen ions and hydroxide ions according to Equation 1.

When the concentration of H⁺ is equal to the concentration of OH^–, the solution is neutral (pH = 7). When H⁺ ions exceed OH^– ions, the solution is acidic (pH < 7). When OH^– ions exceed H⁺ ions, the solution is basic (pH > 7). For example, an aqueous solution of HCl or H₂SO₄ has a greater concentration of H⁺ ions and is therefore acidic. An aqueous solution of NaOH or NH₄OH has a greater concentration of OH^– ions and is therefore basic.

Salts, on the other hand, may undergo hydrolysis in water to form acidic, basic, or neutral solutions. Hydrolysis of a salt is the reaction of the salt with water or its ions. A salt is an ionic compound containing a cation other than H⁺. The broad range of cations and anions that combine to form salts (e.g., NaNO₂, NH₄I, CuSO₄, NaBr) makes it more difficult to predict whether the resulting salt solution will be acidic, basic or neutral. When a salt dissolves, the ions can be identified as acids or bases.

In a dilute salt solution, a soluble salt dissociates completely into its ions. Thus, a water solution labeled “NaCl” actually contains Na⁺ ions and Cl^– ions (Equation 2). The acid–base properties of a salt such as NaCl are determined by the behavior of its ions. To decide whether a water solution of NaCl is acidic, basic, or neutral, the effect of the Na⁺ and Cl^– ions on the pH of water must be considered. In the case of NaCl, the resulting aqueous solution is neutral with a pH of seven. Both Na⁺ and Cl^– do not cause the solution to become acidic or basic. While some ions have no effect on the pH of water, some ions can make the solution acidic because they produce H⁺ ions when dissolved in water, and others can make the solution basic because they produce OH^– ions when dissolved in water. One definition of acids and bases is known as Lewis acids and Lewis bases. A Lewis acid is a substance that accepts an electron pair. A Lewis base is a substance that donates an electron pair.

Let’s take a look at sodium acetate, NaC₂H₃O₂. When sodium acetate is added to distilled water, the solution turns basic. Why? The two ions in sodium acetate are sodium, Na⁺, and acetate, C₂H₃O₂^–. As seen earlier, the Na⁺ ion does not affect the pH of the solution. However, the acetate ion does react with the water and the solution turns basic (Equations 3 and 4). Metal ions can also behave as Lewis acids. For example, let’s look at aluminum chloride hexahydrate, AlCl₃•6H₂O. When this compound dissolves, it breaks into Al(H₂O)₆³⁺ and 3Cl^–. The Al(H₂O)₆³⁺ than acts as a Lewis acid (Equation 5 and 6). Distilled or deionized water can become acidic over time. The longer distilled water is exposed to the air, the more CO₂ that can dissolve into the water. If CO₂ dissolves, the water can become acidic for the same reasons we discussed (Equation 7). To remove the carbon dioxide in the water, the distilled water will be boiled and then cooled.

In this lab, different salts will be tested. The salts will be dissolved in water, the pH of the resulting solutions will be measured, and chemical equations will be written.

Experiment Overview

The purpose of this experiment is to view the acid–base properties that occur during the hydrolysis of salts. Writing and understanding the chemical reactions behind the hydrolysis will also be studied.

Materials

Prelab Questions

1. Define hydrolysis.
2. Define Lewis acid and Lewis base.
3. Sodium bicarbonate is added to water. Three drops of universal indicator are added to the solution and the solution turns purple, indicating the solution is basic. Write the balanced net ionic equation to show the formation of OH^–.

Safety Precautions

Ammonium chloride, potassium chloride, sodium carbonate, sodium phosphate and zinc sulfate are slightly toxic by ingestion and are body tissue irritants. Universal indicator solution is an alcohol-based flammable liquid. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Wash hands thoroughly with soap and water before leaving the laboratory. Please follow all laboratory safety guidelines.

Procedure

1. Measure approximately 150 mL of distilled or deionized water into the 250-mL beaker.
2. Place the beaker on a hotplate. Set the hotplate to heat on high. Move onto Step 6 as you wait for the water to boil.
3. Allow the water to boil for 10–15 minutes.
4. Remove the beaker from the hotplate with the hot vessel gripping device and allow it to cool.
5. Cover with Parafilm^® after the beaker has partially cooled. Note: Do not put the Parafilm on immediately after removing the beaker from the heat.
6. While the water is heating, boiling and cooling, set up the lab station for testing.
7. Obtain six clean test tubes and place in the test tube rack. Note: It’s important the test tubes are clean and free of contaminates. If necessary, clean and dry out the test tubes before proceeding.
8. Mark each test tube with the five salts to be tested. The sixth test tube will be distilled water only.
9. Label each weigh boat with the five salts as well.
10. Measure out 0.4 g of each salt. Note: Use a clean spatula for each salt.
11. Transfer the salt to the correct test tube.
12. When the water has cooled, measure 10 mL of distilled water for each test tube. Pour the water in the test tubes.
13. Place 3 drops of universal indicator into each test tube.
14. Using a clean stirrer for each test tube, stir the solutions and record your observations in the data table on the Hydrolysis of Salts Worksheet.
15. Consult your instructor for appropriate disposal procedures.

Student Worksheet PDF

14100_Student1.pdf