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Nikolay [14]
3 years ago
6

How much time is needed to deposit 1.0 g of chromium metal from an aqueous solution of crcl3 using a current of 1.5 a?

Chemistry
2 answers:
PSYCHO15rus [73]3 years ago
8 0
The metal component of the given compound, CrCl3, is chromium. The number of moles per 1 g of chromium is calculated through the equation below,

        n = (1 g Cr)(1 mol Cr/51.996 g Cr)
              n = 0.0192 mol Cr(3 electrons/1 mol Cr) 
                n = 0.0577 e-

Determine the number in charge by multiplying with Faraday's constant,

      C = (0.0577 mol Cr)((1 F/1 mol e-)(96485 C/ 1F)
             C = 5,566.87 C

Then, calculate time by dividing the charge with the current,

     t = 5566.87 C/1.5 A 
     t = 3711.25 minutes
     t = 61.84 hours

<span><em>Answer: 61.84 hours</em></span>


sergiy2304 [10]3 years ago
3 0

Answer:

t=406.2s=6.77min

Explanation:

Hello,

At first, one must compute the CrCl3 moles by using its molecular mass:

n_{CrCl_3}=1gCrCl_3*\frac{1molCrCl_3}{152.35gCrCl_3}=0.00632molCrCl_3

Now, we can apply the following equation to compute the requested time, based on the Faraday's constant, the involved charge and the fed current:

C=n_{CrCl_3}*F\\C=0.00632molCrCl_3*96485\frac{C}{molCrCl_3}\\C=609.3C\\C=I*t\\t=\frac{C}{I}=\frac{609.3C}{1.5C/s} \\t=406.2s=6.77min

Best regards.

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