In a certain electrolysis experiment involving $\mathrm{Al}^{3+}$ ions, $54,9 \mathrm{~g}$ of $\mathrm{Al}$ is recovered when a current of $0.352 \mathrm{~A}$ is used. How many minutes did the electrolysis last? ( $1 \mathrm{C}=1 \mathrm{~A}-\mathrm{s}$ ) Round your answer to 3 significant digits.
Note: Reference the Fundamental constants table for additional information.
minutes.

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∻Solution∻ ‖ 1 ⋮ **Calculate moles of Al**: First, we need to calculate the moles of aluminum (Al) that were deposited during the electrolysis. The molar mass of Al is $26.98 \mathrm{~g/mol}$. Using the mass of Al recovered, we can find the moles using the formula: $ ext{moles} = \frac{ ext{mass}}{ ext{molar mass}}$. ‖ ‖ 2 ⋮ **Determine total charge**: Each mole of $\mathrm{Al}^{3+}$ requires 3 moles of electrons to be reduced to Al metal. The charge of one mole of electrons (Faraday's constant) is approximately $96500 \mathrm{~C/mol}$. The total charge (Q) needed for the deposition can be calculated by multiplying the moles of Al by 3 and then by Faraday's constant. ‖ ‖ 3 ⋮ **Calculate time**: The charge (Q) can also be expressed as the product of current (I) and time (t), as per the formula $Q = I \cdot t$. By rearranging the formula to solve for time, we get $t = \frac{Q}{I}$. Using the given current and the total charge calculated in the previous step, we can find the time in seconds. ‖ ‖ 4 ⋮ **Convert to minutes**: To find the duration in minutes, we divide the time in seconds by 60. ‖ ∻Answer∻ ⚹ [Insert final answer here] ⚹ ∻Key Concept∻ ⚹ Electrolysis time calculation ⚹ ∻Explanation∻ ⚹ The duration of electrolysis can be calculated by determining the total charge required for the deposition of a metal and then using the relationship between charge, current, and time. ⚹◊What is the relationship between the amount of substance deposited during electrolysis, the current used, and the time taken, as expressed by the formula $Q = I * t$?⍭ Generate me a similar question◊

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