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Electrochemical Conventions

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    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)

    Electrochemical Conventions is shared under a CC BY-NC-SA 4.0 license and was authored, remixed, and/or curated by LibreTexts.

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