It took 150 . s for a current of 1.25 A to plate out 0.109 g of a metal from a solution contai

step 1 So let's assume that the metal does have a plus one charge and calculate its molar mass. We'll t

It took 150 . s for a current of 1.25 A to plate out 0.109...

It took $150 . \mathrm{s}$ for a current of $1.25 \mathrm{~A}$ to plate out $0.109 \mathrm{~g}$ of $\mathrm{a}$ metal from a solution containing its cations. Show that it is not possible for the cations to have a charge of $1+$.

Key Concepts

Valency and Oxidation State in Electrochemistry

The oxidation state or valency of ions defines the number of electrons involved in their reduction or oxidation during electrolysis. By comparing the calculated moles of electrons from the charge passed with the moles of metal deposited, one can validate or rule out a proposed oxidation state. In this problem, determining that the ratio of electrons to metal atoms doesn't match what is expected for a mono-positive ion demonstrates that the cations cannot have a 1+ charge.

Stoichiometry of Electrolytic Reactions

The stoichiometry in electrolytic reactions describes the relation between the electrons transferred and the number of moles of ions reduced or oxidized. When a metal ion is reduced to its elemental form, the number of electrons required per ion corresponds to the ion’s charge. This relationship is essential for checking consistency between the measured mass deposited and the expected amount from the charge passed, particularly when testing a hypothesis about the ion’s charge.

Faraday’s Laws of Electrolysis

Faraday’s laws relate the amount of material deposited at an electrode to the total electric charge passed through the electrolyte. The first law states that the mass of a substance altered at an electrode during electrolysis is proportional to the quantity of electric charge used. This principle is foundational in linking electrochemical measurements to the moles of substance deposited, regardless of the specific substance or reaction details.

Charge and Moles of Electrons

Calculating the total charge (using Q = current × time) and converting this charge into moles of electrons (by dividing by Faraday’s constant) is crucial in electrochemistry. This concept ties together electrical parameters with chemical transformations, allowing the determination of the number of electrons involved in a reaction, which is necessary for assessing reaction stoichiometry and ion valency.

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