Ask Me Help Desk - Gravimetric Analysis URGENT!

Objective: Determine the amount of sulphate by the gravimetric method.

Procedure
A) To precipitate BaSO4

I)Pipette 25.0ml of given sulphate solution (Na2SO4) into a 250ml beaker.
ii) Add 50ml of water and 5drops of concentrated HCl
iii) Heat to boiling and, with vigorous sitrring, add dropwise from a measuring cylinder 10ml of 10% BaCl2 solution
iv) Cover the beaker with a watch glass and allow digestion for 20 minutes
v) Test for complete precipitation by adding a afew drops of BaCl2 to the clear supernatant liquid.

Question 1.Erm.. The precipitating agent is BaCl2 right? I was guessing the purpose of HCl here. Is it to acidify the pH to allow easy solubility and thus slower precipitation? But even if it is, why would we want a slower precipitation when we can speed by the process? Is it due to that lesser compounds would precipitate at a lower pH, thus reducing error in coprecipitation?

Question 2.And also, my lecturer made us skip the test for complete precipitation part. Part AV). I would like to ask how do I determine whether it's a complete or incomplete precipitation. I have a theory that if its an incomplete precipitation, cloudy forms will appear while if it's a complete precipitation, nothing will happen. Is that why we have to make sure BaCl2 is in excess?

Question 3. I realised there were afew errors to the experiment. It did not instruct us to preweigh the watch glass first. I realised there were some condensation occurring on the watch glass. Do we have to preweigh it and if so, why do we have to do this? It would not affect the precipitate right?

Part B) Washing and filtration of Ba2SO4 precipitate
I)Decant the clear supernatant by filtration at the vacuum pump, using a pre weighed crucible.
ii) Wash and swirl the precipitate with about 20ml of warm deionised water.
iii)Use a 'rubber-policeman' to dislodge any particles on the beaker.
iv)Allow to settle
v)Decant the liquid through filter paper
vi)Repeat washing and decanting at least twice.
vii)Discard filtrate

Question 1) Why was warm deionised water used? Won't it cause the solubility of the precipitate to increase,allowing more of it to escape the filter? I felt that iced-cold water should be used to wash the precipitate to maintain its solid form.

Question 2) What is the purpose of letting it settle after swirling it? Is it due to the colloidal form? ( I don't really know what colloidal form means, read it through the internet)

Part C) Drying and weighing of BaSO4 precipitate
I) Dry the crucible in the oven at 150 degree celsius for about 1/2 hour.
ii)Cool the crucible with BaSO4 precipitate in a desiccator for about 10 minutes.
iii) Weigh the crucible when it is cooled down.

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Question 1.Erm.. The precipitating agent is BaCl2 right? I was guessing the purpose of HCl here. Is it to acidify the pH to allow easy solubility and thus slower precipitation? But even if it is, why would we want a slower precipitation when we can speed by the process? Is it due to that lesser compounds would precipitate at a lower pH, thus reducing error in coprecipitation?

The precipitating agent is BaCl2. The purpose of the HCl is to acidify the solution. Barium sulfate (BaSO4) is extremely insoluble. Barium carbonate, BaCO3, will precipitate in basic solution and you don't want that because it would cause an error in the analysis. Carbon dioxide dissolves easily (very easily) in water and most water contains a fair amount of CO2 just by sitting around. If you acidify the solution, you prevent carbonates from precipitating. You'll force the following equilibrium to the left and prevent the formation of carbonate:

$H_2O + CO_2 \rightleftharpoons H_2CO_3 \rightleftharpoons 2H^+ + CO_3^{-2}$

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Question 2.And also, my lecturer made us skip the test for complete precipitation part. Part AV). I would like to ask how do i determine whether it's a complete or incomplete precipitation. I have a theory that if its an incomplete precipitation, cloudy forms will appear while if it's a complete precipitation, nothing will happen. Is that why we have to make sure BaCl2 is in excess?

The lecturer probably made you skip the test for complete precipitation because of time constraints. If Barium is present in excess, the following equilibrium will be forced to the right.

$Ba^{+2} + SO_4^{-2} \rightleftharpoons BaSO_4$

To test for complete precipitation, once the Barium Sulfate has settled you add a drop of additional Barium Chloride. If any cloudiness appears (it will appear where the drop fell), precipitation is not complete.

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Question 3. I realised there were a few errors to the experiment. It did not instruct us to preweigh the watch glass first. I realised there were some condensation occurring on the watch glass. Do we have to preweigh it and if so, why do we have to do this? It would not affect the precipitate right?

The only reason I can think of to preweigh the watch glass is in case some of the mixture spattered and left some material (either sodium sulfate or barium sulfate) on the watch glass. You would weigh the watch glass when it was dry. Afterwards, you would dry the watch glass in an oven and make sure that it weighed the same. If I were doing the analysis, I would simply pick up the watch glass with a pair of clean tweasers, and rinse it off into the beaker (or flask) with a wash bottle.

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Question 1) Why was warm deionised water used? Won't it cause the solubility of the precipitate to increase,allowing more of it to escape the filter? I felt that iced-cold water should be used to wash the precipitate to maintain its solid form.

Yes, it might cause the solubility of barium sulfate to increase, but BaSO4 is so insoluble that you'll never be able to measure the amount that dissolves no matter how accurate your balance is. Warm water increases the solubility of the other salts (sodium chloride and barium chloride) and makes it easier to wash it out. If you don't wash the other salts out, it will cause an error in weighing and the resultant analysis will be wrong.

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Question 2) What is the purpose of letting it settle after swirling it? Is it due to the colloidal form? ( I don't really know what colloidal form means, read it through the internet)

A colloid is a very fine precipitate that doesn't fall to the bottom of the flask -- it stays suspended. Milk is a colloid.

This isn't the reason you wash and swirl it, however. You want to wash and swirl it so as to make sure that any salt, NaCl and BaCl2, is washed from the surface of the barium sulfate. The whole idea is to make sure the precipitate contains no contaminants. You let it settle so that the supernatant can be poured off (decanted) without losing any BaSO4.

Hey. Thanks a lot! =) Now I have a better understanding of what I was doing on that day. So if its in colloid form, it will filtrate through the crucible is that right?

If it were colloidal, you'd have an awful time trying to get it to precipitate. It most likely would go through the crucible -- at least some of it would. You can boil colloids and sometimes get the particles to agglomerate.

ic.. haha. Thanks a lot! =) I sincerely appreciate your help. Kudos to u =)

1) Why do you add HCl to the mixture of BaCl2 and NaSO4?
2) What are the factors that favour the formation of BaSO4 crystals?
3) How do you check the purity of BaSO4 crystals formed?
4) How do you improve the purity of BaSO4?
5) Besides, determining the amount of sulphate by the gravimetric method, what other objectives does this experiments poses for students?
6) What is the theory behind using instruments such as crucible, vacuum pump and desiccator

iii) Heat to boiling and, with vigorous stirring, add dropwise from a measuring cylinder 10ml of 10% BaCl2 solution

For the above statement, is there a reason why we have to carry out vigorous stirring while adding the BaCl2 solution?

Thanks

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Originally Posted by Askl

iii) Heat to boiling and, with vigorous stirring, add dropwise from a measuring cylinder 10ml of 10% BaCl2 solution

For the above statement, is there a reason why we have to carry out vigorous stirring while adding the BaCl2 solution??

Thanks

Its to prevent a relative supersaturation point of the precipitating agent. That is also the reason why the precipitating agent was added in dropwised.Therefore, when the precipitating reagent mixes with the solution, it avoids a local high concentration of the precipitating reagent, avoiding nucleation and colloid formation..

High relative supersaturation would normally occur in this experiment, leading to affecting the particle size as the rate of nucleation is much greater than the rate of growth causing the size of precipitate to be small. This is also known as colloids, causing barium sulphate particles to be lost in the filtrate during filtration.

Q is the concentration of the solute in the solution at the moment the reagents are mixed. S is the concentration of the solute in a saturated solution
Nucleation is the formation of tiny particles around which the bulk of the precipitate grows, leading to particle growth. Nucleation and particle growth competes for molecules/ions that is being precipitated. Therefore, if nucleation is faster than particle growth, a large number of tiny fine precipitate forms, causing colloidal suspension. Some of the precipitate would be lost in the filtrate this way as the pore of the filter paper is bigger than the particle size. But if particle growth is faster than nucleation, large particles would be formed.

Therefore, the precipitation was done from a very dilute solution to keep Q small. When Q is small, the relative supersaturation would be low.

The precipitating reagent was also added in dropwise, and stirred vigorously to keep Q small, preventing a local high concentration of the precipitating agent.

When the precipitating reagent mixes with the solution, it avoids a local high concentration of the precipitating reagent, avoiding nucleation and colloid formation.

Precipitation was also done in a hot and acidic solution to increase the amount of solute that can be in solution at equilibrium.

Treatments can be done to remove the colloidal form. The using of a volatile salt can be used to wash the precipitate, displacing ions of opposite charge that are attracted to the nucleus known as counter ions. This would then produce a denser precipitate that is easier to filter. The volatile salt can be removed during drying of the precipitate.
This is what I stated in the final report that I did.

Okay, thanks so much! You're a great help. :)

Why do we need to use deionized water in dissolving the unknown mass? I know that one reason would be so that the ions in water do not react with the substance being measured, but what is another reason?

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Why do we need to use deionized water in dissolving the unknown mass? I know that one reason would be so that the ions in water do not react with the substance being measured, but what is another reason?

That's the primary reason - contamination. Your reagents must be pure, so the water you dissolve things in must also be pure. You don't want to precipitate something that isn't what you think is the product.

Raw water might contain suspended solids which you can't see with the naked eye which might cause you to err.