Question

A constant current of 0.350 A is passed through an electrolytic cell containing molten CrCl₂ for 21.7 h. What mass of Cr(s) is produced? The molar mass of chromium is 52.0 g/mol. Provide your answer rounded to 3 significant digits.​

Answer 1

An electrochemical cell can generate or use electrical energy. The mass of solid chromium that will be deposited on the electrochemical plate is 7.17 gm.

What is current?

Current in an electrochemical cell is the ratio of the quantity of electricity in columns and time in seconds.

Given,

Current (I) = 0.350 A

Time = 21.7 hours

Molar mass of chromium = 52.0 g/mol

First time is converted into seconds:

1 hour = 3600 seconds

21.7 hours = 76020 seconds

The quantity of electricity flowing in the electrochemical solution is calculated as:

$\begin{aligned} \rm Q & = \rm It\\\\& = 0.350 \times 76020 \\\\& = 26607\;\rm C \end{aligned}$

Electricity required for depositing 1 mole or 52.0 g chromium is calculated as:

In electrochemical solution, chromium chloride is dissociated as:

$\rm CrCl_{2} \rightarrow Cr^{2+} + 2 Cl^{-} \\\\\rm Cr^{2+} +2 e^{-} \rightarrow Cr$

Two moles of electrons are needed to deposit 52.0 g of chromium.

If, 1 electron = 96500 C

Then, 2 electron = 193000 C

The mass of chromium deposited is calculated as:

193000 C = 52 g chromium

So, 26607 C = $\dfrac{26607 \times 52}{193000} = 7.17 \;\rm gm$

Therefore, 7.17 gm of chromium is produced.

Learn more about an electrochemical cell here:

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