Assume that the volume of each solution in Figure 20-22 is 100.0 mL. The cell is operated as an

step 1 Hello, so today we're going to be looking at this cell right here. And at first we're going to b

Assume that the volume of each solution in Figure 20-22 is...

Key Concepts

Stoichiometry of Redox Reactions

Stoichiometry in redox reactions involves balancing the equations for oxidation and reduction processes and deducing the relationships between the amounts of electrons transferred and the moles of reactants/products formed. It is essential for determining how changes in the quantities of species, due to electrolysis, impact the cell’s performance when it subsequently operates as a voltaic cell.

Nernst Equation

The Nernst equation provides a way to calculate the cell potential under non-standard conditions by accounting for the effect of ion concentrations on the electrode potentials. It relates the measured cell potential to the standard cell potential, temperature, number of electrons transferred, and the reaction quotient, making it crucial for understanding how concentration changes during electrolysis affect the spontaneous voltage of a voltaic cell.

Electrolytic versus Voltaic Cells

An electrolytic cell uses electrical energy to drive a non-spontaneous reaction, whereas a voltaic (or galvanic) cell harnesses the energy released by a spontaneous redox reaction to generate electrical power. Understanding the transition from an electrolytic to a voltaic mode is important, as it involves changes in the cell’s reactant concentrations that influence the cell potential.

Faraday's Laws of Electrolysis

Faraday's laws relate the quantity of electric charge passed through an electrolyte to the amount of substance that undergoes oxidation or reduction at an electrode. These laws are used to calculate the number of moles of electrons transferred during an electrolytic process, taking into account the current and the duration for which the current is applied.

Electrolysis

Electrolysis is the process of driving non?spontaneous chemical reactions using electrical energy. In this process, an external power source forces electrons to move through an electrolyte, leading to oxidation at the anode and reduction at the cathode. The amount of material transformed during electrolysis can be quantified in terms of the charge passed through the system.

Transcript

00:01 Hello, so today we're going to be looking at this cell right here.

00:05 And at first we're going to be operating this cell as an electrolytic cell.

00:11 And we're going to be running 0 .5 amperes through this cell for 10 hours.

00:20 So basically, this would be the zinc right here.

00:25 We have this zinc electrode with one molarity of zinc on.

00:31 Ions in the solution and this is going to be the cathode since this is an electrolytic cell and the zinc is going to be reduced and over here is going to be our anode since this is the electrolytic cell and we're going to have copper being oxidized so usually it would go in the reverse but we're running electricity through this so that this is the reaction that takes place.

01:31 So let's first figure out how many moles of electrons we're going to be moving.

01:42 So an amp here is one kulum per second.

01:50 And since we have 10 hours, let's convert the hours into minutes.

02:04 And then minutes in the seconds and then we have 0 .5 coulamps per second and for every coulom we have you can just use fairaday's constant for every 96 ,485 coulamps we have one mole of electrons moving around so if we do some math we will see that we have 0 .187 moles of electrons moving around.

03:32 So let's figure out, so we know how much moles of electrons are moving around, but what exactly does that mean? well, the zinc will takes two moles of electrons to be reduced and copper produces two moles of electrons for every copper, one mole of copper that's converted to an ion.

04:02 So first, let's see how much zinc ions we have.

04:08 So we have 100 milliliters and one molarity, which is one mole over one liter.

04:17 So that means that we have 0 .1 .0 moles of zinc ions.

04:31 And we have 0 .10 moles of copper ions.

04:43 So for every 0 .187 moles of electrons, we see that zinc is being reduced, so we're removing zinc ions from the solution.

04:59 So if we have 0 .187 moles of electrons, then for every 2 moles of electrons is, 1 mole zinc 2 plus reduced or removed from solution.

05:27 So we see that we are removing 0 .093 3 moles of zinc 2 plus.

05:45 And then we see conversely for the moles of electrons that we have moving around.

05:57 For every two moles of electrons, one mole of copper will be produced.

06:28 So we are removing 0 .0933 moles of zinc and we are adding 0 .0933 moles of copper.

06:43 So now let's say that we stop running electricity through this cell and it becomes a electrolytic cell.

06:58 I mean it becomes a voltaic cell...

Please give Ace some feedback