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Chemistry LibreTexts

Electrochemical Conventions

I.  General Oxidation-Reduction (Redox)

  1. Oxidation = loss of electrons.  (Losing electrons increases the charge (oxid. state).)
  2. Reduction = gain of electrons.  (Gaining electrons reduces the charge (oxid. state).)

Can add LEO says GER and OIL RIG etc.

  1. Oxidizing Agent:
  • causes an oxidation
  • takes electrons
  • is itself reduced
  1. Reducing Agent:
  • causes a reduction
  • gives electrons
  • is itself oxidized

II.  Electrochemistry

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.”


  • Fred Wood (UC Davis)