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

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 dialogue-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.

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