An Overhead Equilibrium

Introduction

A colorless solution becomes dark orange upon addition of a solution and then a solid. The dark orange color disappears after the addition of another solid but reappears again when more solution and the original solid are added.

Concepts

• Complex ions
• Equilibrium
• Le Chatelier’s principle

Materials

Safety Precautions

Potassium thiocyanate is toxic by ingestion and emits toxic fumes if strongly heated. Iron(III) nitrate solution may be skin/tissue irritant. Sodium phosphate monobasic (monohydrate) is moderately toxic by ingestion. Wear chemical-resistant gloves and chemical splash goggles. 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. The resulting solution may be flushed down the drain with excess water according to Flinn Suggested Disposal Method #26b. Clean up spills of iron(III) nitrate solution immediately because iron(III) nitrate solution easily stains.

Procedure

1. Using a 50-mL graduated cylinder, measure out 20 mL of potassium thiocyanate solution. Transfer the potassium thiocyanate solution to a Petri dish, and place the Petri dish on an overhead projector. Turn on the overhead projector.
2. Add 5 drops of iron(III) nitrate solution in different spots in the Petri dish. Note that the orange spots produced are a little darker than the iron(III) nitrate solution.
3. Swirl the solution until the orange color is uniform throughout.
4. Add ½ pea size amount of the potassium thiocyanate crystals in one spot. A dark orange spot results. Wait about 30 seconds so the students can observe the movement of the dark orange color through the solution.
5. Swirl the solution to dissolve the crystals and the dark orange color will become uniform throughout.
6. Add ¼ pea size amount of the sodium phosphate monobasic crystals in one spot. Wait about 60 seconds as the color in the vicinity of the crystals becomes lighter than the rest of the solution. This is a great example of ion diffusion in a solution.
7. Swirl the solution to dissolve the crystals and until the solution is colorless throughout.
8. Add one drop of the iron(III) nitrate solution in one spot off to the side. Note the blood red color. Don’t stir.
9. Add a pea size amount of the potassium thiocyanate crystals in a different spot. Wait about 30 seconds so the students can observe that the area around the crystals becomes orange.

Teacher Tips

• Enough materials are included in this kit to perform this demonstration seven times; one practice, five classes and one for “Murphy’s Law.”
• Steps 8 and 9 were incorporated into the demo based on the question: “Are the Fe³⁺ and SCN^– ions still present even though there is no color?”
• Put the chemical equations on the overhead while you are doing the demonstration. You can explain or, better yet, have the students explain, the equilibrium shifts.
• Beral pipets or wood splints are ideal to stir the solution.

Sample Data

Answers to Questions

1. Write chemical equations for the following reactions.
   1. The reaction between Fe³⁺ and SCN^– ions

      Fe³⁺(aq) + SCN^–(aq) → FeSCN²⁺(aq)

   2. The reactions between Fe³⁺ and H₂PO₄^– ions

      Fe³⁺(aq) + H₂PO₄^–(aq) → [FeH₂PO₄]₂₊

2. What direction does equilibrium for the first reaction above shift when ferric nitrate or potassium thiocyanate is added?

   Equilibrium in the reaction shifts to the right.

3. Why did the solution turn colorless when the sodium phosphate monobasic was added?

   The addition of the sodium phosphate monobasic shifts equilibrium to the left because of the competing reaction between the Fe³⁺ and H₂PO₄^– ions. The product of that reaction is colorless.

Discussion

Fe³⁺ and SCN^– ions form the complex ion FeSCN²⁺, which is dark red in color (Equation 1). Addition of Fe(NO₃)₃ or KSCN increases the concentration of these reactants and causes the equilibrium in Equation 1 to shift in the forward reaction, or to the right. This observation illustrates the effect of LeChâtelier’s Principle—a change in any reaction condition causes the equilibrium to shift in a direction that counteracts the effect of the change.

Addition of NaH₂PO₄•H₂O results in the equilibrium shifting in the reverse reaction (to the left). This is due to a competing reaction of Fe³⁺ and H₂PO₄^– ions to form a colorless [FeH₂PO₄]²⁺ complex. This effect can be reversed again by adding more Fe³⁺ or SCN^–.

References

Special thanks to Jim and Julie Ealy, The Peddie School, Hightstown, NJ, for providing us with the instructions for this activity.