Chemistry 30

Reversible Reactions: Acid-Base Indicators - Lesson 1.1

Acid-base indicators are a quick and easy way to demonstrate reversible systems. Here are a few examples using different indicators:

• Add a few drops of phenolphthalein to a test tube half-filled with distilled water.  To this colourless solution add a few drops of 0.1 M NaOH. The solution will turn light pink. Adding additional base will cause the colour to deepen.

Next add several drops of 0.1 M HCl. Swirl the test tube to mix. The solution will turn colourless.

Phenolphthalein Colour	pH range
colourless	0 - 8
pink	8 - 12
colourless	12 and above

• Add about 1 mL of thymol blue to 100 mL of distilled water in an Erlenmeyer flask, noting the colour. Divide this into two test tubes, using one as a control. Thymol blue in distilled water is blue.

Add a single drop of 0.1 M HCl to the test tube that is not the control. Swirl the tube to ensure complete mixing. Continue adding HCl until a definite colour change is observed and have students record the new colour.

Continue adding more 0.1 M HCl until a second colour change occurs. Record.

Reverse the colour change by the addition of 0.1 M NaOH.  As with the HCl, continue adding NaOH until two colour changes have been noted.

Thymol Blue Colour	pH range
blue	9 and higher
green	8.5
yellow	2 - 8
red	0 - 2

Reversible Reactions - Gases - Lesson 1.1

In Lesson 1.1 the following equilibrium is described:	2 NO2 (g)		N2O4 (g)
	Red		Colourless

USE CAUTION - concentrated acid

PROCEDURE

1. Place about 10 mL of concentrated nitric acid in a flask.
2. Carefully add a penny. A deep red gas, NO₂, will form.
3. Carefully "pour" the gas into two test tubes and stopper them.
4. Place one tube in a beaker of boiling water. The colour gets a deeper brown.
5. Place the other tube in an ice bath. The gas will become almost colourless.
6. Return both tubes to room temperature. The gases will return to their original colour.

The Reaction:   

Cu_(s) + 4 H⁺_(aq) + 2 NO₃^-_(aq) → Cu²⁺_(aq) + 2 NO_{2 (g)} + 2 H₂O_(l)

Le Châtalier's Principle - Equilibrium involving Copper(II) ions - Lesson 3.1

• PROCEDURE
  
  Place approximately 2 mL of 0.1 M CuSO₄ in a test tube. CuSO₄ is a light blue solution.
  
  Add several drops of 1 M NH₃. The result will be a light blue precipitate. The addition of more NH₃ causes the solution to turn a darker blue.
  
  Next add several drops of 1 M HCl until a change is noted - the solution will become a lighter blue.

• EXPLANATION
  
  An ammonia solution (Reaction 1) produces ammonium and hydroxide ions. When added to the CuSO₄ solution, OH^– reacts with the Cu²⁺ to produce the precipitate, Cu(OH)₂ (Reaction 2).

• Adding H⁺ (as HCl) removes OH^– from the solution (Reaction 1). This shifts the equilibrium in Reaction 1 towards the right (the products), resulting in a decrease in [NH₃].  This reduction in [NH₃] causes the equilibrium in Reaction 3 to shift to the left (the reactants)

Reaction 1:
NH3 (aq) + H2O(l) NH4+ + OH–(aq)
Reaction 2:
Cu2+(aq) + OH–(aq) Cu(OH)2 (s)
Reaction 3:
	Cu(H2O)42+(aq) + 4NH3 (aq)		Cu(NH3)42+(aq) + 4H2O(l)
	light blue		dark blue

Le Châtalier's Principle - Equilibrium involving Cobalt (II) ions - Lesson 3.1

This easy-to-do demonstration can be used to illustrate both the effect of concentration and temperature on equilibrium systems. The cobalt chloride solution can be kept and reused for future classes.

View a demonstration

• MATERIALS

0.2 M Cobalt chloride hexahydrate, CoCl₂ · 6H₂O

Prepare a 0.2 M cobalt chloride solution by dissolving 2.6 g CoCl₂ · 6H₂O in 100 mL of distilled water.

concentrated (12M) HCl - caution!
hot water bath
ice water bath

• SAFETY
  
  Use extreme caution when handling the hydrochloric acid.
  

• PROCEDURE
  
  The cobalt equilibrium system is represented by the following equation:
  

Co(H2O)62+(aq) + 4 Cl-(aq)		CoCl42-(aq) + 6 H20(l)	ΔH = +50 kJ
pink		blue
hydrated cobaltous ion		chloro complex

Part 1. The Effect of Concentration

1. Begin with 50 mL or so of the pink hydrated form of the cobalt ion in a larger (250 mL) beaker or Erlenmeyer flask.
   
   
2. Carefully add concentrated HCl to the solution until the blue form of the equilibrium system appears.
   
   
3. Add distilled water to the solution in the beaker or flask. The equilibrium will shift back to the pink form.
   

Part 2. The Effect of Temperature

1. Add some pink form of the system to a test tube and place it into the hot water bath, or carefully warm the solution using a bunsen burner. The solution will turn blue.
   
   
2. Place the blue test tube in the ice water bath; it will return to its pink colour.

• EXPLANATION
  
  Part 1. The effect of concentration. When HCl is added to the pink hydrated form of the cobalt ion, the concentration of Cl^- increases. (H⁺ remains in solution as a spectator ion). The increase in Cl^- causes the equilibrium to shift to the product side, and the solution turns blue.
  
  Adding water will shift the equilibrium to the left, forming more pink hydrated cobaltous ion.
  
  Part 2. The effect of temperature. As written above, the forward reaction is an endothermic reaction. According to Le Châtalier's Principle adding heat favours the endothermic direction. Thus warming the system will favor the blue chloro complex side of the reaction. When placed in an ice bath, the exothermic direction will now be favored, shifting the equilibrium to the pink hydrated side.