15. Electrode potential, galvanic cell
Electrode reaction, standard hydrogen electrode, Nernst-equation, electrode potential scale, gas electrode, metal electrode, galvanic cell/battery, oxidation/reduction ability
Electrochemistry:
-(1785) Martinus van Marum electrolyzed zinc and antimony by an electrostatic generator
-(1799) Volta column → real electrolysis possibility
-(1800) William Nicholson and Anthony Carlisle used the Volta column to decompose water
-(1807) Humprey Davy made the discovery that instead of K, Na water electrolyzes
-Davy’s student Michael Faraday discovered oxidation degree, anode, cathode, electrode and ions (scientists thought the atom is the smallest particle)
Electrode potential
Potential: The capability of something
Electrode potential: the ability to make the charged particles move
Electrode reaction
-in the Zn → Zn2+ +2e- half equation the formation of Zn2+ is favored (in comparison with what?)
-but where does the electron go? charge separation is unfavored
-in the Cu2+ + 2e- → Cu half equation the formation of Cu is favored (in comparison with what?)
-but where does the electron come from? charge separation is also unfavored here
-The answer to the previous question is that we cannot measure the electrode potential of a single electrode, but the difference of the two electrode’s potentials can be measured
-The difference is the output voltage (U) of the galvanic cell or can also be called the electromotive force (EMF) of the galvanic cell, or the potential of the galvanic cell (Ecell)
-In order to measure the potential difference between the two electrodes we need to make a galvanic cell, which consists of the two electrodes and a diaphragm/salt bridge which allows the transfer of charges:
-The other thing we need to measure the difference between the potential of the two electrodes is a reference
-The reference is the Standard Hydrogen Electrode (SHE)
-The Standard Hydrogen Electrode is an electrode made from platinum
-the platinum is surrounded by hydrogen gas at atmospheric pressure and 25°C
-The electrode is submerged in a solution with a pH of 0 (hydrogen ion concentration is 1 mol/l)
-The corresponding half-reaction is: H+ + e- = ½ H2
-The potential of the SHE is 0 by definition
-The output voltage of a galvanic cell made from any electrode and the SHE gives the (standard) electrode potential of the given electrode (can be +/-)
Nernst-equation
-we can express what influences the electrode potential with this equation
-we can see that the electrode potential is dependant on 5 things
-the standard potential of the electrode (it is the measured voltage under standard conditions (pH=0 → c=1 M, atmospheric pressure, 25°C), when a cell is connected to a SHE)
-the temperature
-the amount of electrons in the half-equations
-the concentration of the oxidized form (the product of the anode)
-the concentration of the reduced form (the product of the cathode)
-(R (8,314 j/mol*K) and F (96485 c) are constant)
If the temperature of the galvanic cell is 25°C we rewrite the equation like this:
For metal electrodes
-the reaction is heterogeneous (the reduced form is a solid metal, while the solution is a liquid so we have two phases)
-since the reduced form is a metal the concentration of it doesn’t need to be taken into account in the Nernst-equation:
For gas electrodes:
-the concentration of the gas can be can be calculated as the pressure of the gas over the atmospheric pressure (c=p/p0 where p0 = atmospheric pressure)
-so their potential is pressure dependent
-we can calculate the pH using gas electrodes from their electrode potential
-for hydrogen:
-for chlorine:
Second order electrode
-we put a silver rod (electrode) into a silver-nitrate solution saturated with silver-chloride precipitate
-Ksp=[Ag+][Cl-] → [Ag+]=Ksp/[Cl-]
-Half-reaction: Ag=Ag+ +e-
-Nernst for the half reaction:
εAg/Ag+=ε0Ag/Ag++0,059lg[Ag+]
-if we substitute we get:
εAg/Ag+=ε0Ag/Ag++0,059lg Ksp/[Cl-]
-using the properties of logarithm:
εAg/Ag+=ε0Ag/Ag++0,059lg Ksp-0,059lg[Cl-]
-since lg Ksp is constant we can incorporate it into ε0Ag/AgCl
ε0Ag/AgCl= ε0Ag/Ag++0,059lg Ksp
-with this we can rewrite the equation as:
εAg/AgCl= ε0Ag/AgCl -0,059[Cl-]
Calculating K (equilibrium constant) using the voltage output of the galvanic cell (Electromotive force)
-in equilibrium EMF/E =0 →the galvanic cell is exhausted
-the standard potentials determine the equilibrium:
-the agent with the more positive (HIGHER) standard potential is the oxidizing agent (the one that gets REDUCED on the CATHODE)
-the agent with the more negative (LOWER) standard potential is the reducing agent (the one that gets OXIDIZED on the ANODE)
-proof of this:
Electrode potential scale
-ε can only decrease if the concentration of he oxidized form decreases
-potetnial is an intense state funcion
-it must equilibrate, the larger one decreases, the lower one increases by the concentrations of the oxidized and reduced forms
-there is a scale for this (where the system’s potential is equilibrated):
Galvanic cell/battery (reversible systems)
-a galvanic cell converts chemical energy into electric energy, while the battery is discharged
-properties of a galvanic cell/battery:
-energy density (electrode potential, charge)
-reliability
-safety
-charging time
-lifetime
-capacity
-specific capacity/mass
.chemical energy storage
Reduction and oxidation ability
-if a metal reacts with metal ions one of them will oxidize (lose electrons), one of them will reduce (take up electrons) according to the reaction (this also happens when a halogenide reacts with an other halogenide ion)
-according to standard electrode potentials, the one with the lower is oxidized, the one with the higher is reduced
-if the two assumptions match, the reaction happens, if not it doesn’t, this is important in active corrosion protection