ChemCollective-Virtual-Labs/C3/Determination-of-Equilibrium-constant/English-timed

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Time Narration
00:01 Welcome to this tutorial on Determination of Equilibrium Constant using Vlabs.
00:08 In this tutorial, we will learn, To determine equilibrium constant for Cobalt chloride reaction.
00:15 Observe the effect of change in temperature and concentration on equilibrium.
00:22 To follow this tutorial, you should be familiar with, ChemCollective Vlabs interface.
00:29 If not for relevant tutorials please visit our website.
00:34 Here I am using,

Mac OS version 10.10.5

00:39 ChemCollective virtual labs version 2.1.03

Java version 8.

00:48 Here I have opened Vlabs interface.
00:52 Click on File menu. Scroll down to Load Homework option.
00:57 Default Lab Setup dialog-box opens.
01:02 From the list, double-click on Chemical Equilibrium.

Two options appear.

01:10 Double-click on Cobalt Lab option.
01:14 Stockroom Explorer has required solutions and Problem Description.
01:20 Double-click on Problem Description.

Problem description states that, we need to apply Le Chatelier's principle, for aqueous Cobalt chloride reaction.

01:33 Using the equilibrium concentration we need to find, equilibrium constant for cobalt chloride reaction and

Effect of temperature and reactant concentration on equilibrium.

01:47 Let us define chemical equilibrium.
01:50 Chemical equilibrium is a state of the reversible reaction when two opposing reactions occur at the same rate.
01:59 Concentration of the reactants and products do not change with time at equilibrium.
02:06 This side shows a general Equilibrium Reaction.
02:11 This is the equation for Equilibrium Constant.
02:15 The chemical equation for this reaction is shown here.
02:20 Cobalt forms complexes with water molecules, as well as chloride ions.
02:27 A solution of Hexaaquacobalt(II)complex is pink.
02:33 When hydrochloric acid is added to the solution, the colour changes to blue.
02:39 This corresponds to the formation of Cobalt Chloride complex.
02:44 Equilibrium Constant changes with, Change in concentration of reactants or products and
02:51 Change in temperature.
02:54 Click on Workbench.
02:57 Double-click on Cobalt chloride experiment solutions cabinet.
03:03 Double-Click on 1M Cobalt Chloride.
03:07 A flask with 100 mL of 1M Cobalt Chloride is added to the workbench.

Observe the colour of the solution in the flask.

03:18 Solution Info panel shows the required information.
03:23 Colour of the solution is pink due to presence of Hexaaquacobalt(II)complex.
03:30 Note the concentrations of Hexaaquacobalt(II)complex, chloride ions and Cobalt chloride in your observation book.
03:40 Double-Click on 12 M HydroCloric acid.
03:44 Click on Glassware menu. Select Erlenmeyers.
03:50 From the list click on 250 mL Erlenmeyer flask.

Rename the flask as A.

03:59 From the Glassware menu, select Pipets.
04:03 From the list, click on 25 mL Pipet.

From the Glassware menu, select 50 mL Buret.

04:12 Using Pipet, measure 25 mL of Cobalt chloride solution.
04:18 Click on Withdraw.

Keep the flask aside.

04:23 Place the Pipet over flask A.

Type 25 and Click on Pour.

04:31 Keep the Pipet aside.
04:34 Fill the buret with 50 ml of 12 M hydrochloric acid.
04:40 Bring 12 M hydrochloric acid flask on to 50 mL buret.
04:46 Type 50 in the Transfer amount input bar.

Click on Pour.

04:52 Keep the flask aside.
04:55 Bring the buret on to flask A.
04:58 Add hydrochloric acid from the burette in 1 mL increments, using Precise Transfer mode.
05:07 Type 1 in the Transfer amount input bar and click on Pour.
05:13 Similarly transfer another 6 mL of hydrochloric acid using Precise Transfer mode.
05:22 We have transferred 7 mL of hydrochloric acid to Flask A.
05:28 Notice the change in temperature on the thermometer.
05:33 Temperature decreases during the reaction.

It means that, the reaction is endothermic.

05:41 The total volume of solution in Flask A shows 32 mL.
05:47 Observe the colour change in Flask A, colour changes to brown.
05:53 It may take a few seconds to reach equilibrium state.

Now the concentrations of the reactants and products are constant.

06:03 Note the values of the concentrations of Hexaaquacobalt(II)complex, chloride ions and cobalt chloride in your observation book.
06:14 Continue the titration.

Pour 1 mL at a time into the flask.

06:21 Transfer another 8 mL of Hydrochloric acid to flask A.
06:27 Now we have transferred 15 mL of hydrochloric acid to flask A.

Total volume of liquid in flask A is 40 mL

06:40 Note the colour change in Flask A.

Please wait for the reaction to reach equilibrium condition.

06:49 Again note the concentrations of Hexaaquacobalt(II)complex, chloride ions and cobalt chloride in your observation book.
07:00 Similarly add another 3 ml of hydrochloric acid to flask A.

Total volume in flask A is 43 mL.

07:11 Note the colour change in flask A.

Note the concentrations in your observation book.

07:19 Finally add 5 mL of hydrochloric acid from the buret.

Total volume in flask A is now 48 mL.

07:30 Note that we have added 23 mL of hydrochloric acid to flask A.
07:37 Colour of the solution in flask A is blue.
07:41 Again note the concentrations of Hexaaquacobalt(II)complex, chloride ions and cobalt chloride at equilibrium in your observation book.
07:52 Let us see how to calculate Equilibrium Constant.
07:56 Calculate Equilibrium Constant using the given formula.
08:01 Substitute the values of concentrations of cobalt chloride, Hexaaquacobalt(II)complex and chloride ions in the equation.
08:11 This is the value of equilibrium constant after pouring 7 mL of hydrochloric acid.
08:19 Similarly, here are the values of equilibrium constant for 15, 18 and 23 mL of hydrochloric acid.
08:29 Since, temperature is constant, equilibrium constant values are almost the same.
08:36 Switch to workbench.

Next I will demonstrate the effect of temperature on equilibrium.

08:45 Earlier we have observed that, this reaction is endothermic.

It means, heat is absorbed during the reaction.

08:54 Le Chatelier’s Principle states that,

If an equilibrium is disturbed by changing the conditions, position of equilibrium moves to counteract the change.

09:06 According to the principle, for endothermic reactions, rate of forward reaction increases with increase in temperature.
09:14 Back to workbench. Keep the burette aside.
09:20 Let us increase the temperature of the reaction flask to 35 degree Celcius.
09:26 Right-click on flask A, from the context menu, select Thermal Properties.

Input box opens.

09:35 In Set the temperature to text box, type 35.

Check the box for Insulated from surroundings.

Click on OK.

09:46 Thermometer shows 35 degree Celcius.
09:50 Note the values of concentrations of Hexaaquacobalt(II)complex, Chloride ions and Cobalt Chloride.
09:59 This is the calculated value of equilibrium constant at 35 degree Celcius.
10:05 Compare it to equilibrium constant value with that at 25 degree Celcius.
10:12 Note that the value of Kc at 35 degree Celcius is greater than Kc value at 25 degree Celcius.
10:21 This is because the reaction is an endothermic reaction.
10:26 As the temperature increases rate of forward reaction increases.

Hence more product is formed.

10:34 Switch to workbench,

Remove the used pipette and burette from the Workbench.

10:42 Right-click on flask A

From the context menu, select Thermal Properties.

10:49 Un-check the box for Insulated from surroundings.

Click OK.

10:55 This will bring the temperature back to 25 degree Celcius.
11:00 Now let us remove the chloride ions from the reaction using Silver Nitrate.
11:07 Double-Click on 6M Silver nitrate from Stockroom Explorer.

From the glassware menu, select 25 mL Pipet.

11:18 Withdraw 25 mL of silver nitrate(AgNO3 ) into the 25 mL Pipet.
11:26 Transfer 25 mL of silver nitrate(AgNO3) from pipet to flask A in 5 mL increments.
11:34 Type 5 in Transfer Amount input bar.

Click on Pour.

11:40 Note the temperature. It increases as you add silver nitrate to flask A.
12:00 Note the colour change.
12:02 It indicates the formation of Hexaaquacobalt(II)complex.

Click on Solid radio button.

12:11 Note the amount of silver chloride(AgCl) in grams column.
12:16 Silver nitrate(AgNO3) reacts with chloride(Cl-) ions in solution to form silver chloride(AgCl).
12:23 Here chloride(Cl-) ions decrease in the solution, to compensate for the deficit,

Rate of reverse reaction increases.

12:33 Cobalt chloride complex decomposes to form HexaaquaCobalt(II) complex.

This example is a proof for LeChatelier's principle.

12:44 Let us summarize.
12:46 In this tutorial, we have learnt, To determine equilibrium constant for Cobalt chloride reaction.
12:53 Observe the effect of change in temperature and concentration on equilibrium.
12:59 As an assignment,

Prepare a solution by adding 25 mL of Cobalt chloride solution and 23 mL of Hydrochloric acid.

13:09 Add 40 mL of water in 10 mL increments to the prepared solution

Observe the colour in the flask.

13:19 Calculate Equilibrium Constant before and after addition of water.
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14:00 This tutorial is contributed by Snehalatha kaliappan and Madhuri Ganapathi from IIT Bombay.

Thank you for joining.

Contributors and Content Editors

PoojaMoolya