18.4 Standard Electrode Potentials Maggie Hanson

We just learned that the standard cell potential for an electrochemical cell depends on the specific half-reactions occurring in the half-cells and is a measure of the potential energy difference between the two cells Each electrode has its own potential too, called the standard electrode potential The standard cell potential is just the difference between the two standard electrode potentials

Water as an analogy for standard electrode potential Water spontaneously flows from high to low potential energy When two half-cells are connected, electrons flow from the electrode with more negative charge (higher potential energy) to the electrode with more positive charge (less potential energy) Low Potential Energy High Potential Energy DIRECTION OF SPONTANEOUS FLOW

Continuing the Water Analogy The water analogy is useful but we run into a problem… We can measure the water level in a tank, but we can’t measure the electrode potential in a half cell directly. We can only measure the overall potential that occurs when the two half-cells are combined in a whole cell. Here’s the solution: We arbitrarily assign a potential of 0 to the electrode in a particular type of half-cell and then measure all other electrode potentials relative to that zero

We normally choose the standard hydrogen electrode (SHE) to have a potential of 0: An inert platinum electrode immersed in 1M HCl with hydrogen gas at 1atm bubbling through the solution When SHE acts as the cathode, the following half-reaction occurs: 2H + (aq) + 2e - H 2 (g) E o cathode = 0.00V If we connect the SHE to an electrode in another half cell, we can measure the difference in the potential between the two electrodes. Since we know the potential of the SHE is 0, we now have the potential of the other electrode

E o cell = E o final - E o initial = E o cathode - E o anode We define E o cell as the difference in voltage between the cathode (final state) and the anode (initial state), since electrons travel from anode to cathode The measured cell potential (E o cell ) for this cell is +76 volts. So, we can find the E o anode using this formula -76V In this case, the potential for the Zn/Zn+ half- cell is negative, indicating that an electron at the Zn/Zn+ electrode has greater potential energy than it does at SHE (more negative voltage=higher potential energy because negative charge repels electrons)

If we connected an electrode in which the electron has more positive potential to the SHE, then that electrode would have a more positive voltage (lower potential energy) (so the opposite of what we just saw). In this case, electrons would NOT spontaneously flow from anode to cathode If we did this using Cu/Cu 2+ as the anode, we would find that the E o anode is +.34 V

Main ideas in standard electrode potentials: 1. The electrode potential of the standard hydrogen electrode SHE is exactly zero 2. The electrode in any half-cell with a greater tendency to undergo reduction is positively charged relative to the SHE and therefore has a positive E o a. The electrode in any half-cell with a lesser tendency to undergo reduction (or greater tendency to undergo oxidation) is negatively charged relative to the SHE and therefore has a negative E o 3. The cell potential of any electrochemical cell (E o cell ) is the difference between the electrode potentials of the anode and the cathode 4. E o cell is positive for spontaneous reactions and negative for nonspontaneous reactions

Example Use tabulated standard electrode potentials to calculate the standard cell potential for the following reaction (the equation is balanced): Al(s) + NO 3 - (aq) + 4H + (aq) Al 3+ (aq) + NO(g) + 2H 2 O(l) The standard electrode potential for Al/Al 3+ is -1.66V The standard electrode potential for NO 3 - /NO is.96V Answer on next slide

E o cell = E o cat - E o an =.96V - (-1.66V) =2.62V Is the reaction spontaneous?

Predicting the Spontaneous Direction of a Redox Rxn 1. Look at the electrode potentials for each half-reaction 2. The reaction with the more negative electrode potential tends to lose electrons and therefore undergo oxidation 3. The reaction with the more positive electrode potential tends to gain electrons and therefore undergo reduction

Consider the two half-reactions: Ni 2+ (aq) + 2e - Ni(s) E o = -.23V Mn 2+ (aq) + 2e - Mn(s) E o =-1.18V The manganese half-rxn has a more negative electrode potential, so it will repel electrons and proceed in the reverse direction (oxidation) Nickel thus proceeds in the forward direction To confirm this, treat manganese as the anode and nickel as the cathode, then plug the E o values into the formula and make sure E o is positive (the reaction is spontaneous)

Example Predict whether the following redox reaction is spontaneous. If the reaction is not spontaneous, write an equation for the spontaneous direction in which the reaction would occur. Fe(s) + Mg 2+ (aq) Fe 2+ (aq) + Mg(s) E o Mg 2+ /Mg: -2.37V E o Fe 2+ /Fe: -.45V Solution on next slide

Not spontaneous! The magnesium half-reaction has the more negative electrode potential and therefore repels electrons more strongly and undergoes oxidation, while iron undergoes reduction. As written, the reaction is not spontaneous. However, the reverse reaction is spontaneous. *another way you can tell this is nonspontaneous is by using E o cell = E o cat - E o an. E o is negative since the cathode is more negative than the anode.

Resources All of the standard electrode potentials are on page 873 in the textbook Good review of redox reactions and electrochemistry Crash course electrochemistry Learning objectives 3.12 and 3.13! See description.pdfhttp://media.collegeboard.com/digitalServices/pdf/ap/ap-chemistry-course-and-exam- description.pdf