I. General Oxidation-Reduction (Redox)
- Oxidation = loss of electrons. (Losing electrons increases the charge (oxid. state).)
- Reduction = gain of electrons. (Gaining electrons reduces the charge (oxid. state).)
Can add LEO says GER and OIL RIG etc.
- Oxidizing Agent:
- causes an oxidation
- takes electrons
- is itself reduced
- Reducing Agent:
- causes a reduction
- gives electrons
- is itself oxidized
a) Electrodes = Where the oxidation/reduction (redox) reactions occur
- Anode = Where oxidation half reaction occurs. (begin with vowels).
- Cathode = Where reduction half reaction occurs. (begin with consonants).
b) In the external circuit (ie, the wire) the electrons always flow from the oxidation site to the reduction site.
- Therefore, electrons in the external circuit always flow from the anode to the cathode.
c) Voltaic (Galvanic) Cells- an electrochemical cell in which a spontaneous reaction produces electricity.
- The cell potential is always positive. (If you got a problem where the voltage was negative, it means the cell is spontaneous in the reverse direction.)
- Electrons move in the external circuit from the negative electrode to the positive electrode.
- Thus, in the voltaic cell, anode = negative and cathode = positive
d) Electrolytic Cells- an electrochemical cell in which a non-spontaneous reaction is carried out by electrolysis.
- (Electrolysis: the decomposition of a substance (in a molten state or in an electrolytic solution) by an electrical current.) The voltage is always negative and a metal is plated out or a gas is evolved.
- Electrons are forced to move in the external circuit from the positive electrode to the negative electrode.
- Thus, in the electrolytic cell, anode = positive and cathode = negative
e) Drawing cells (both voltaic and electrolytic)
The negative electrode (not necessarily the anode) is shown on the LEFT I don’t think this is necessarily true.
Thus, the left-hand electrode is the:
- anode if voltaic cell
- cathode if electrolytic cell
f) Line Diagrams (Cell Diagrams)
- the left hand electrode is the anode (where oxidation occurs)
- The sign on the left hand electrode is negative
- a boundary between different phases is represented by a single line (÷).
- reactants are at the left of the double line (||) which represents the boundary between the half cells (usually a salt bridge); products are at the right of |.
- Different species in the same solution half cell compartment are separated by commas. | Pt(s) ê Cl2(g) êCl-(aq) || Pb2+(aq), H+(aq) ê PbO2(s)|
- (some professors do not care about the order of the ions presented in a line diagram, while others say that, ions can be ordered in each half cell so that increasingly more positive ions are nearest the double lines. Example for a voltaic cell:
anode electrode(s) ÷ A-1 (aq, B+1(aq) ÷÷ C+2(aq), D-1(aq)÷ cathode electrode(s))
g) Standard Reduction Potentials (or voltages)
- voltages of a half cell in which all gases are 1 atm and all solutions are 1 M (25oC)
- half cells are listed as reduction processes.
- the more positive the voltage, the more easily reduced.
- thus, the strongest oxidizing agents are found at the top left (as F2 (g)) and the strongest reducing agents are found at the bottom right (as Li(s))
- a useful memory aid is “upper left reacts with lower right.”