14. Electrodes and Potentiometry
Potentiometry - Redox Reactions
Potentiometry - Line Notation
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So here we're going to say that potentially alma tree uses electrodes to measure voltages that provide vital chemical information on solutions. Now we're going to say here, the difference in potential between the two electrodes provide an analysis for the composition of the an elite. Now the measure of the voltage and electrode relative to the reference is done in the absence of current. So there's no current that's being run through this instrument here, we're gonna say that are non standard cell potential sl equals the potential of our indicator electrode minus our reference electrode plus the cell potential of our junction potential. But here the junction potential comes from both ends of the salt bridge and this occurs because there's maybe an imbalance in the concentration of ions from one side to the other as the salt bridges are connected to the solutions of the cathode and an ode. Or there is a difference in the size of ions within the salt bridge and the concentrations in those two have cell reactions. Now we can say here with proper use and proper ions in the salt bridge, we can minimize the junction potential self potential to the point where it's negligible. So here we're going to say that our reference electrode again is E R E F your salt bridges. What's causing this cell potential or junction potential to occur if you use the incorrect ions. Normally we want we want to use potassium chloride as the best choice of ionic inert ions within your salt bridge because the ions are relatively the same size. Then we have our analytics solution and then we have our indicator electrode. Now here we are given a basic example which would classify as having our reference electrode in our indicator electrode here on the left side we have our an ode in the anad compartment. We have our silver solid with our chloride ion. They undergo an oxidation. So we've lost an electron from the silver ion making it plus one which causes it to combine with the chloride ion as a precipitator. Solid here, we've lost our electrons. So here are the electrons as a product. So electrons are leaving and going towards the cathode here we have our platinum wire. The platinum wire here is just an inert electrode that's being used within the solution so that electrons can basically go onto its surface and crust themselves onto the surface. Here we have our iron three ions absorbing an electron to become iron two. Plus now we're going to say when it comes to the reference electrode versus the indicator electrode, we're gonna say the reference electrode is indicative of the anodes and in the reference electrode, this is part of the cell that is held constant. So its concentration isn't changing. And then here the indicator electrode, it's part of the cell that contains the an elite being analyzed. Now this portion that in this box can also be compartmentalized into just this small portion here. So they're the same exact reaction. Except in the one on the left, we're putting them in separate jars. But both processes are the same. Now we'd say that we know that the half reactions in both of them are given as this. So here we can see that this is our cathode. This here is our a note. We can see here because your cathode has the larger cell potential and you're an it has a smaller cell potential. That this is a spontaneous process. Because if we did cathode minus an ode through our calculations we get a cell potential that's greater than zero, which would tell us that we're dealing with the galvanic or voltaic cell and therefore it is spontaneous. Now here the nurse equation provides a mathematical relationship between the electrodes potential and and analyze reduced and oxidized forms. So here we're gonna say the difference in potential is based on only one of the half cell concentrations because remember your reference electrode, its concentration is not changing its staying constant. So it's not really what's going to determine our difference in potential, it's the varying concentrations of our reference of our indicator electrode that will do that here. If we take a look on this side we have our cathode and here we have our anodes. So for our cathode compartment we'd say that our self potential for cathode under nonstandard conditions equals the self potential of our cathode under standard conditions minus 0.5916 divided by the number of electrons transferred here. This would be log. And remember here it's products. Overreact Ints products. This concentration overreact ints Here. We only have one mole of electrons being transferred. We plug in these values. The concentrations here to give us our whatever value would be for our cathode under nonstandard conditions. Then here for a node, it's the same exact equation. Except now the concentration of chloride ions is not being affected, it's staying the same. And here it's also products Overreactions but we ignore solids and liquids. So we get this value here. Once we have both values, we could then figure out what our overall cell potential would be by doing cathode minus an out. Which is the same thing as saying here, indicator minus your reference. And that's how we get ourself potential here. So remember here now we're using our our potential that we can calculate to determine the concentration of a particular an elite within a given solution. So that's the whole point. Shows us a connection between the varying concentrations and the direct impact that will have on our voltage. Will it cause my overall cell potential to increase or decrease with changes in those concentrations
Potentiometry uses electrodes to measure voltages that also provide vital chemical information on their solutions.