What a great video . I'm chemistry edu student from Indonesia. Thanks for the explanation sir. I really need this. I'd like to use your video as a reference for my final project. Hope you don't mind about it. Thank you. Best regards ~ Sriza Hayati
Thank you for this video! I am definietly doing this experiment with my class! I do have a question about the hydrated Co complex. Did you buy the Co and then mixed it with water? Or did you buy the cobalt hydrated complex online? Hope you anwser this question!
Because the Ag+ will combine with the Cl- to form a precipitate (form a solid) thus the Cl- concentration is decreased. The stress of the decreased Cl- ion will cause the reaction to "shift" to the left to increases the Cl- concentration to relive the stress. A precipitate that occurs effectively takes the free chloride ions out of the solution and locks them into the ionic crystal lattice.
@@colephillips8070 Yes I do think that adding the solution does provide more water which will contribute BUT the precipitate being made is a larger factor in driving reverse reaction.
What if KCl was added instead of HCl? Will the solution turn blue as well? I did this experiment last week, but there wasn't any change in the colour. It only turned blue when heated. Why is it so? :/
Saturated KCl solution is only about 4.6 M compared to 12 M HCl. If your solutions are made using water as the solvent the equilibrium will already be shifted far to the left since water is a stronger ligand than cl- A shift right (Blue) will take a lot of cl ions so we need the 12 M HCl. However, if you make the solution with ethanol like in video it will be easier for the cl- ligands to form. So I think KCl would do the trick
What was the chemical name of the precipitate that formed from the double replacement reaction in question 8? And is the heat on the left side of the equation?
You cannot include water in your equilibrium equation that determines the equilibrium constant because water has no molarity or concentration. The law of mass action requires the components of the equation that determine the equilibrium constant have a concentration (molarity of a solution) or partial pressure if the chemicals are in the gas phase. Chemicals that are in the solid phase or liquid phase do not have concentrations while aqueous solutions and gases can. Gases can be considered a solution and can be given a concentration (molarity) as well as a partial pressure. The reason for this is that we track the the equilibrium constant through changes of entropy and solutions can have a change in entropy as they dilute or become more concentrated while liquids (like water) and solids do not.
MrGrodskiChemistry thanks for your reply. I struggle to explain this one. So when adding water is it true to use Le Chateliers principle by saying an increase in the amount of water - not concentration - drives the reaction to the left? Or is it more about the decrease in concentration having a greater effect on the left side thus the shift to the left? That would seem to relate to the equilibrium equation which ignores the water.
@@mattyoxide3650 If you are teaching the concept of Le Chatelier's Principe there is nothing wrong with explaining the increase or decrease of water to drive the reaction BUT the reason that amount water has this effect will require a deeper understanding of equilibrium. As mentioned before, water does not have a molarity but the other components of the chemical equation does. So you are correct in that the water amount will influence the concentration of one of the complex ions more than the other thus driving the reaction. Increasing the amount of water will decreases the concentration of all of the ions on both sides of the equation but the decrease will be less with the cobalt chloride complex on the right side . If you look at the reaction quotient, Q, which has the same equation as the equilibrium constant Keq we will find the answer. The 2 products are on the denominator and one has a power of 4 (Cl-) thus decreasing all the concentrations of all ions the same will cause the denominator to be smaller than the numerator thus increasing the reaction quotient larger than the equilibrium constant (which will only change due to temperature). If the Q is larger than the Keq than the reverse reaction will be more spontaneous and there will be a shift to the left when water is added.
@@jrmcr1458 I know that is the answer the instructor is looking for but is it also due to the addition of water found in the Silver Nitrate solution? After all the addition of water shifted equilibrium to the pink side
Yes. NaCl has the same common ion as Cl ion in the reaction. It would have the same effect as any chemical species that would provide Cl ions in solution.
@@MrGrodskiChemistry thank you for your concern and suggestion. I'm a chemistry teacher in Israel. Cobalt is permitted to use in solutions. of course all necessary cautions will be taken.
Well the reaction is the one written above the demonstration. The reaction shifts or becomes more spontaneous in one direction as you add this salt. I do not want to totally explain this here as students are using this video for lab. Please email me and I will explain further if you desire.
The preparation of the complexes is purely qualitative however I would advise against this preparation due to the toxicity and hazards that Co presents.
Thanks for posting this! I needed a make-up for my students and the only video I could find showed the Heat in the equations. I also want them to determine the thermochemistry.
Students are using this lab as an assignment and I would rather not that answer that directly. The answer to that question is presented with the observation in the cold/hot water test. If you know whether an exothermic or endothermic reaction is favored when heated or cooled, you have the answer.
