An equilibrium mixture in a 1.00L flask of the above treaction contains the following gases: 0.13M of I2, 0.26M of F2, 0.21M IF5, and 0.38M I4F2. To this mixture, some iodine was added. When equilibrium reestablished, 7.6g of F2 was found to be present. What mass of I2 was added to the system?
Submitted by equilibrium2014 on Sun, 2014-10-05 18:42
I performed a lab for the following equilibrium but there were some questions I could really use some help on explaining events in the lab through Le Chateliers Principle.
The Equilibrium Equation: CoCl4^-2 + 6H20 <=> (Co(H20)6)^+2 + (4Cl)^-1
1. I placed the test tube of the equilibrium solution in a hot water bath, followed by a cold water bath. The original blue solution became a darker blue in the hot water (shifted left) and it turned a lighter blue in the cold water (shifted right). As heat is exiting in this reaction I said it was an exothermic reaction. I am asked to explain my observations using Le Chatelier's Principle but I'm not too sure how I should exactly do that.
The Principle: An equilibrium system subjected to a stress will shift to partially alleviate the stress and restore equilibrium.
2. I then had to add some AgNO3 (Silver Nitrate) to the equilibrium solution. It became pink. How do I explain this too using Le Chatelier's Principle? How did this affect the concentration changes of the solution?
3. I then added some CaCl2 (solid) to the equilibrium before turning blue. How do I explain the colour change in relation to Le Chatelier's Principle? How did this affect the concentration changes of the solution?
4. I then added some water. It turned from a blue colour to a pink and then almost clear. How do I explain this color change using Le Chatelier's Principle? How did this affect the concentration changes of the solution?
Is this question actually solvable? It was on a quiz I just took and I couldn't figure it out...
Equation is 2I2+2H2S <-> 4HI+2S2. Given: KC=2.33x10-5, 0.250 mol HI in a 2L flask. Same with S2. Want: Equilibrium concentration of I2.
So I set up the ice table after putting everything in molarity.
2I2 + 2H2S <-> 4HI + S2
0 0 .125 .125
+2x +2x -4x -x
2x 2x .125-4x .125-x
Plugged it into ratios:
[HI]4[S2] / [I2]2[H2S]2 = ((.125-4x)4(.125-x)) / (16x4) = 2.33x10-5 I assumed the 4x and x in the numerator would be small so that changed it to (.1255)/(16x4). Plugging and chugging, I end up with a value of somewhere near 0.54 = x. So 2x= 1.06. Now, that just didn't look right, but I couldn't figure how else to do this. Any help?
Ammonia decomposes at high temperatures. In an experiment to explore this behavior, 3.00 moles of gaseous NH3 were sealed in a rigid 1.50 L vessel. The vessel is heated at 800K and some of the NH3 decomposes in the following reaction: NH3 (g) ---> N2 (g) + 3H2 (g)
The system eventually reaches equilibrium and is to contain 2.18 moles of NH3. What are the values of Kc and Kp for this reaction at 700K.
A solution was prepared such that the initial concentrations of Cu+2 and CN- were 0.0120 M and 0.0400 M, respectively. these ions react according to the following equation - Cu2+ + 4CN- > Cd(CN)4 ^2-. Kc = 1.0 x 10^25. What will be the concentration of CN- at equilibrium?
I made an ICE table and plugged everything into Kc,
so 1.0 x 10^25 = x / (0.0120 - x)((.04-4x)^4), is that right so far?
Submitted by Spinnerphoenix on Fri, 2014-04-18 11:29
Mixtures of propan-2-ol and propanone can be spearated by distllation due to their different boiling points. Explain why these compounds have such different boiling points even though they have very similar molar masses.
when we deal with the raoult's law equation for an ideal solution of 2 ideal liquids which are volatile, do we have to use the number of moles at equilibrium to calculate the mole fraction of solvent in the solution? or can we use the initial number of moles ?