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Chemistry Forums for Students => Undergraduate General Chemistry Forum => Topic started by: iop388 on March 28, 2020, 07:24:45 PM

Title: Aluminium Metal Electrolysis
Post by: iop388 on March 28, 2020, 07:24:45 PM
How many seconds are required to produce 12.00 g of aluminum metal from the electrolysis of molten Al2(SO4)3 with an electrical current of 4.50 A?
Title: Re: Aluminium Metal Electrolysis
Post by: billnotgatez on March 28, 2020, 08:14:08 PM
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Title: Re: Aluminium Metal Electrolysis
Post by: AWK on March 29, 2020, 03:35:58 AM
The problem could be calculated if aluminum sulfate were stable at the melting point, but this is not the case.
Title: Re: Aluminium Metal Electrolysis
Post by: Borek on March 29, 2020, 04:08:30 AM
Quote from: AWK on March 29, 2020, 03:35:58 AM
The problem could be calculated if aluminum sulfate were stable at the melting point, but this is not the case.

I don't see how the decomposition changes amount of Al produced in the electrode process?
Title: Re: Aluminium Metal Electrolysis
Post by: AWK on March 29, 2020, 04:18:49 AM
Quote from: Borek on March 29, 2020, 04:08:30 AM
Quote from: AWK on March 29, 2020, 03:35:58 AM
The problem could be calculated if aluminum sulfate were stable at the melting point, but this is not the case.

I don't see how the decomposition changes amount of Al produced in the electrode process?
Melting point of Al2O3 is much, much higher
Title: Re: Aluminium Metal Electrolysis
Post by: Borek on March 29, 2020, 04:58:20 AM
It makes the process unrealistic, still doesn't change the applicability of the Faraday's law.
Title: Re: Aluminium Metal Electrolysis
Post by: AWK on March 29, 2020, 05:11:12 AM
Quote from: Borek on March 29, 2020, 04:58:20 AM
It makes the process unrealistic, still doesn't change the applicability of the Faraday's law.

I formally agree with your statement but it should be clearly stated in the task, as for example in gas tasks it is advisable to treat real gas as an ideal gas.

Students remember all the nonsense best.