[Rutherford]

Reactions of complex formation are frequently used in titrimetric methods of determination of various inorganic ions. For example, fluoride forms a stable complex with aluminum(III):
6F^-+Al³⁺ = AlF₆^3–
In water the complex gives a neutral solution. This process can be used for the direct titration of fluoride and indirect determinations of other species.

In the first experiment, a sample solution containing fluoride was neutralized with methyl red, solid NaCl was added to saturation, and the solution was heated to 70–80°C. The titration was performed with 0.15 M AlCl₃ until yellow color of the indicator turned pink.

1.What process occurred at the endpoint?
2.Why heating increased the endpoint sharpness?
3.What is the purpose of adding sodium chloride?


I think that the process is:
6F^-+Al³⁺ = AlF₆^3–
I don't know 2 and 3. Maybe heating increases the reaction rate. Any help would be appreciated.

[Borek]

I think that the process is:
6F^-+Al³⁺ = AlF₆^3–

That's not what the question asks - you are expected to explain the color change.

[Rutherford]

The pink color would mean acidic environment. The complex is neutral so what could cause the color change?

[Hunter2]

Think what happens if all Fluoride is consumed to the aluminum chloride.

[Rutherford]

AlCl₃+3F^- AlCl₃F₃^3-?

[Borek]

The pink color would mean acidic environment. The complex is neutral so what could cause the color change?

These are both very important points. Actually they contain the key to the answer.

[Rutherford]

Maybe:
AlCl₃+H₂O Al(OH)₃+H⁺+Cl^-
Acidic H⁺ ions are produced here, but Al makes a precipitate so it can't react with the chlorine. I am really out of ideas  .

[Hunter2]

Check here: http://www.rod.beavon.clara.net/AlCl3_and_water.htm

[Rutherford]

So a gas is produced.

Then I have [Al(H₂O)₆]³⁺ in the solution. Then it hydrolyses to produce the precipitate (Al hydroxide). How would it react with fluoride?

[Borek]

Check here: http://www.rod.beavon.clara.net/AlCl3_and_water.htm

Nonsense - that would work in the case of ANHYDROUS AlCl₃. You have SOLUTION of AlCl₃.

List all substances present in the neutral solution.

[Rutherford]

Then what's actually happening here  ?

[Hunter2]

Check here: http://www.rod.beavon.clara.net/AlCl3_and_water.htm

Nonsense - that would work in the case of ANHYDROUS AlCl₃. You have SOLUTION of AlCl₃.

List all substances present in the neutral solution.

And this is written in the second part of the article. Please read all and dont call nonsense. Aluminiumsalts in water are acidic. Did you ever shave your bear wet. If you get cut you can stopp the bleeding with aluminiumsulfate, because of acidic reaction with the blood.

To solve the miracle Aluminium get complexed by flouride until it is consumed, the free aluminium gives acidic reaction, what can be seen by change of colour of the indicator.

[Rutherford]

Thanks.
Isn't that what I wrote here:

I think that the process is:
6F^-+Al³⁺ = AlF₆^3–
I don't know 2 and 3. Maybe heating increases the reaction rate. Any help would be appreciated.

Then the rest of Al³⁺ hydrolyses to produce the hydroxyde as you said. What about question 2 and 3?

[Borek]

And this is written in the second part of the article. Please read all and dont call nonsense. Aluminiumsalts in water are acidic.

Sorry, my comment was addressed at Raderford, who blindly claimed that the gas is produced.

To solve the miracle Aluminium get complexed by flouride until it is consumed, the free aluminium gives acidic reaction, what can be seen by change of colour of the indicator.

There is another process that can be responsible for the pH change. Solution we start with after neutralization contains some HF, and Al³⁺ grabs all F^- it can. That shifts the HF dissociation equilibrium. Judging which of these processes is dominating is not easy without exact calculations.

[Rutherford]

Never mind then, I skip this.