Explaining the Result of Adding AgNO3 to Test Tube #6

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Discussion Overview

The discussion revolves around the chemical reaction that occurs when a small amount of AgNO3 is added to a solution containing CoCl4. Participants explore the resulting color change and precipitate formation, seeking to understand the underlying chemical processes involved.

Discussion Character

  • Homework-related
  • Exploratory
  • Debate/contested

Main Points Raised

  • One participant notes that the solution turned pink after adding AgNO3, interpreting this as a shift towards the products and observing the formation of a precipitate.
  • Another participant proposes a new equilibrium equation involving AgNO3 and questions how the equilibrium can be shifted to the right.
  • Concerns are raised about charge balance in the proposed equations, suggesting that more Co(H2O)6^2+ may need to be produced to explain the color change.
  • Some participants challenge the correctness of the reactions written, specifically questioning the interactions between Ag+ and Cl- in the solution.
  • A later reply suggests that AgNO3 is well soluble and questions the necessity of using "HOH" instead of "H2O" in the equations.

Areas of Agreement / Disagreement

Participants express disagreement regarding the correctness of the chemical equations presented. There is no consensus on the accurate representation of the reactions occurring in the solution.

Contextual Notes

Participants have not reached a resolution on the correct chemical equations or the implications of the observed color change and precipitate formation. There are also unresolved issues regarding charge balance in the proposed reactions.

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Homework Statement



COCL4 + 6H2 <=> Co(H20) +4CL

blue <=> pink

Question: Add a very small amount of AgNO3 to test tube #6. Stopper the test tube and shake.

The solution in test tube turned pink, so I interpreted that as a shift to the products. Also, a precipitate (proper term? a solid was present...) formed.



The Attempt at a Solution




Now I have to offer an explanation for why this occurred. It seems as though the Ag+ is slightly soluble in Cl-, and the Cl is found on the product side, so was the Cl consumed somehow and thus produced CoCl42-?

To me that doesn't seem reasonable.

EDIT: I somehow thought myself into a circle, because CoCl42- wouldn't have been produced, for if it had the solution should have turned blue right? I'm fairly sure my explanation should hinge on the fact that Ag+is only slightly soluble with Cl-, but I don't want to go off on a tangent either, so if I'm horribly misguided please advise.
 
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I'm using this thread as a sort of running journal. If that is unacceptable, for some reason, please do inform me.

New equation:

2AgNO3(s) + 6H2O(al) + 2CoCl4(al)2- <=> Co(h2O)6(al)2+ + 8Cl(al)- + Co(NO3)2 (al) + 2Ag (s) + energy

How many different ways can the equilibrium be shifted to the right? i.e. Removal of product, addition of reactant and removal of energy.
 
The charges on the new equation don't balance. Looks like more Co(H20)6 (al) 2+ will have to be produced, and if that holds true then I can explain the colour change.
 
None of the reaction you have written so far is correct.

What happens in the solution containing Ag+ and Cl-?
 
Borek said:
None of the reaction you have written so far is correct.

What happens in the solution containing Ag+ and Cl-?


CoCl42- + 6HOH + 4AgNO3(s) <=> Co(HOH)62+ + 4AgCl(s) + 4NO3- + energy

Is this equation correct?
 
Much better, but not necesarilly correct.

AgNO3 is well soluble. And no need to put some fancy looking dihydrogen monoxide HOH in the equation :wink:
 
Doesn't HOH = H2O? I just found it easier to type HOH, so I did. Thanks for the help.
 

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