Equilibrium Lab

Equilibrium

Just this week I'm reviewing equilibrium with my IB Chemistry seniors after they finished some summer study on the topic. One of our classes was spent manipulating a classic equilibrium involving copper ions and a copper-chloride complex ion.

Cu2+(aq)      +  4 Cl(aq)        ⇌       CuCl42—(aq)

Blue          Colorless                         Yellow

The lab handout is attached below in both Word and PDF format. (Note: Hat tip to Mr. Tony Locke, who shared this lab with an IB workshop I attended. I have modified it some along the way.) The lab is a pretty basic look at causing disruptions to an equilibrium and making predictions - and then observations - based on Le Châtellier's Principle. I don't think the lab is all that groundbreaking; many of you probably have a similar lab you've used over the years. But I still like it for its simplicity and easy comparisons to the original equilibrium mixture.

Of all the test tubes, though, the students have the most trouble with the results from test tube 5. In this system, the students simply dilute the sample with water and observe any color change. The most common prediction is for a slightly lighter color (but the same general hue). And when students complete this lab and give a quick comparison, it often seems as though there hasn't been much of a color change. But that doesn't reflect the reality in this case, as the equilibrium will shift towards the reactants.

I ran the demo for my class today and recorded the following video:

        Attempt 1 at filming test tube 5 from the equilibrium lab.

Did you see any color change?

It is a barely perceptible shift towards blue.

But at the end of the day, I was putting things away and happened to look down the two test tubes. From this angle the color change was quite noticeable. So I remade my video filming from the top of the test tubes to see if the results were better.

         Attempt 2 at filming the demo for test tube 5 from the equilibrium lab.

Sometimes all that is needed is a different perspective.

I'd love to know what you use for equilibrium, and for challenging student misconceptions. I'm also curious about the language used for equilibrium, such as "shift the equilibrium to the reactants" and other similar phrases. Unless the temperature is changed, the value of Kc doesn't change. Therefore I think this language hinders student understanding of equilibrium and the significance of Kc. Any thoughts on language that can clear this up?

 

Concepts: 

Equilibrium

Le Châtelier's Principle

Transition Metal Complexes

Time required: 

Approximately 45-60 minutes is usually enough time to complete all six test tubes, and to answer most of the questions and have some discussion about the results.

Materials: 

copper II chloride equilibrium mixture (described in Preparation below)

Each group needs the following:

6 test tubes in a test tube rack

small sample (3-5 grams per group per class) of aluminum chloride

dropper bottle of ~1M Pb(NO3)2 solution

access to a hot water bath/hot plate and large beaker to create hot water bath

access to ice to create cold water bath

access to distilled water

Background: 

The following mixture is provided in equilibrium:    Cu2+(aq) + 4 Cl(aq) <-->  CuCl4​2-(aq)

                                                                              Blue        Colorless                Yellow 

Procedure: 
  1. Read through the ENTIRE procedure before moving to the next step.
  2. Observe the color of the equilibrium mixture and record in the data table.
  3. Obtain 6 clean test tubes in a test tube rack and place about 6 cm of the equilibrium mixture in each. (Note: Add extra to test tube 3, but leave test tube 4 empty at this point.)
  4. Record the initial color of each test tube, and the change that will take place. (e.g. Test Tube 1, write “add heat”) Based on your knowledge of equilibrium, predict the color of the test tube AFTER the change. Record in the data table. (Admittedly test tube 1 and 2 are a bit of a guess. For the purpose of making your prediction more meaningful, assume the reaction is exothermic.)
  5. Perform each of the following actions on the various test tubes and record the appropriate results. 

TEST TUBE 1. Warm the mixture GENTLY but do NOT bring to the boil. Check on the test tube after about 10-15 minutes.

TEST TUBE 2.Place the test tube in a beaker of ice water and leave it for about 10 minutes.

TEST TUBE 3.Add some solid aluminum chloride and shake the mixture gently to dissolve the aluminum chloride completely. (Think about how this will change the concentration of the substances involved in the equilibrium.)

TEST TUBE 4.Using some of the solution remaining from Test Tube 3, create Test Tube 4. Then add lead (II) nitrate solution to the mixture. Let the test-tube settle before making observations. This will be a double- replacement reaction, forming lead (II) chloride as the precipitate. (Think again about how this will change the concentration of the substances in the equilibrium mixture.)

TEST TUBE 5.Add an equal volume of distilled water.

TEST TUBE 6.Control test tube. No action required. Use this test tube for color comparisons. 

Observations 

 

 

 

Questions: 

1. Based on test tube 1 and 2, is the forward reaction exothermic or endothermic? Justify your response with data from the lab and sound chemistry principles.

2. Based on the results, predict and explain the magnitude of the equilibrium constant at room temperature. (In other words, is Kc much greater than 1, approximately 1, or much smaller than 1?) Justify your answer.

3. Based on the results, predict and explain the magnitude of the equilibrium constant at room temperature. (In other words, is Kc much greater than 1, approximately 1, or much smaller than 1?) Justify your answer. 

4. State and explain whether the value of Kc increased or decreased when you increased the temperature. 

5. Write the chemical equation, complete ionic equation and net ionic equation for the double-replacement reaction that occurs in test tube 4. 

HL/Extension Questions (related to Topic 13!)

  1. What is the ligand in this reaction?

  2. What is the coordination number of the coordinated complex?

  3. Explain why the complex is colored.

  4. Would a complex with Cu1+ have been colored also? Why or why not? 

Preparation: 

To make the equilibrium mixture I give to students, I use some copper II chloride, and add some sodium chloride or aluminum chloride to shift the equilibrium a bit more towards the products. I don't use a specific concentration, as the lab is very qualitative.

Each group uses maybe 30-50 mL of the equilibrium mixture.

Attribution: 

This lab is modified from the original provided by Mr. Tony Locke, Atlanta International School. 

Join the conversation.

Comments 2

Scott Milam's picture
Scott Milam | Tue, 08/23/2016 - 18:55

for the green copper complex here with concentrated HCl https://www.youtube.com/watch?v=p2qHgoCjRmk and here is an alternative lab using cobalt that shifts between blue and pink with purple intermediately https://www.youtube.com/watch?v=r-cQAF1S4PU. In Lowell's you can also use this as a demonstration for Beer's law where adding water to a colorful mixture that does not change color will keep the color constant when viewed from the top because the light passes through the same amount of chemical (longer cell path, lower concentration) but is lighter in color from the side due to the dilution.

Lowell Thomson's picture
Lowell Thomson | Thu, 08/25/2016 - 07:58

Hi Scott,

Thanks for the comments here. I like the idea of incorporating Beer's Law for a complex that won't change color (and equilibrium position) with dilution to show the effect of concentration and pathlength on absorbance. 

I've been thinking I needed to find some new equilibrium reactions and will explore the two ideas you suggested here.

Thanks.
Lowell