CHAPTER 9 0XIDATION AND REDUCTION
(TOPICS 9 AND 19) SUMMARY
Basic concepts
Oxidation: Loss of Electrons (LEO); Increase in oxidation number (historically: gain of O or
loss of H)
Reduction: Gain of Electrons; Decrease in oxidation number (GER) (historically: gain of H
(or loss of O)
Oxidation Number: Represents a charge on a mono-atomic ion or the apparent charge on an
element if the more electronegative element is assigned the bonding electrons. Quantitatively,
it reflects the oxidation state of an element; an increase in oxidation number by 1 implies a
transfer of one e-: Na(s)Na+ (aq) + e-.
An element has an oxidation number = 0. For a mono-atomic ion, oxidation number of
element = charge on ion. For a multi-atomic ion, sum of oxidation numbers = charge on ion.
Some elements have the same oxidation number in all their compounds (e.g., Na+). Multiple
oxidation states are common for d-block elements; the oxidation number is then written in
Roman numerals to indicate the oxidation state of the element, e.g.: Fe2+: Fe(II); Fe3+: Fe(III).
An oxidizing agent oxidizes another substance; it is reduced in the process (gains electrons);
it experiences a decrease in oxidation number; thus, found in its higher/highest oxidation state:
e.g.: Cr2O72-; MnO4-; F2.
A reducing agent reduces another substance; it is oxidized in the process (loses electrons); it
experiences an increase in oxidation number; thus found in its lower/lowest oxidation state:
e.g., Na, Ti3+, I-.
Reactivity series: A list of elements in order of decreasing reducing ability: An element
higher in the series displaces a lower one from its ions: Cu (s) + 2 Ag+ (aq) 2Ag (s) + Cu2+
(aq). For oxidizing agents, a stronger oxidizing agent displaces a weaker one from its ions: Cl2
(g) + 2 Br- (aq) Br2 (aq) + 2 Cl- (aq).
Electrodes: conductors at which reactions occur in electrochemical and electrolytic cells.
Anode: Anions, negative ions, go to anode; electrode at which oxidation occurs (anodic
oxidation).
Cathode: Cations, positive ions go to cathode; electrode at which reduction occurs (cathodic
reduction).
A cell contains electrodes placed in an electrolyte solution; the electrodes are connected to an
external circuit.
An Electrochemical (Galvanic or Voltaic) Cell uses a spontaneous redox reaction to convert
chemical energy to electrical energy. It consists of two half-cells: Oxidation occurs at the more
negative half-cell. This contains the reducing agent that is oxidized at the (-) electrode, the
anode and gives up electrons to the external circuit. Reduction occurs at the more positive
half-cell (and thus contains the oxidizing agent that is reduced). The salt bridge maintains
electrical neutrality: anions go to the anode, cations go to the cathode. For a current to flow, a
potential difference between the two half-cells must exist this is called the cell potential or
e.m.f. of the cell.
Cell Diagram: The more negative electrode is written first, to the left, followed by the more
positive electrode:
Zn(s)/Zn2+(aq) Ð#Ð#Cu2+(aq)/Cu(s) E°cell = E°red + E°ox = (+0.34) + (+0.76) = + 1.10 V.
© IBID Press 2007 1
CHAPTER 9 0XIDATION AND REDUCTION
(TOPICS 9 AND 19) SUMMARY
Electrode potentials
The Standard Hydrogen Electrode (SHE) is assigned a value of 0 V under standard
conditions of 25°C, solution concentrations of 1.0 mol dm-3 and 1 atmosphere pressure for
gases. Positive E°cell indicates a spontaneous reaction ( "G°); negative E°cell means a non-
spontaneous in the forward direction, or a spontaneous reverse reaction.
Standard electrode/reduction potential is the e.m.f. of a half-cell connected to the standard
hydrogen potential (at 25°C, 1.0 mol dm-3 solution concentrations and gas at 1 atmosphere
pressure). If reduction occurs at the standard half-cell, its sign is +; its sign is negative if
reduction occurs at the standard hydrogen electrode (see Table 15).
By convention, a cell diagram is written so that the oxidation half-reaction metal/metal ions
are placed first, on the left and the reduction half reaction metal ions/metal placed next on the
right of the cell diagram; a Ð# (or / ) is placed between the metal and its ions and the two
aqueous solutions are then separated by a salt bridge represented by two complete or dashed
lines:
Cell Diagram: Zn(s)/Zn2+(aq) Ð#Ð#Cu2+(aq)/Cu(s) E°cell = + 1.10V
E°cell is the difference between the two standard electrode potentials. It is common to list: E°cell
= E°red (right hand side) E°red (left hand side) = + 0.34 V ( 0.76) = +1.10V. It is also possible to
determine the cell voltage from the two half oxidation and reduction reactions by using the
equation:
E°cell = E°red + E°ox.
For the above cell, the two half reactions are:
Zn(s) Zn2+(aq) + 2 e ; E°ox.= +0.76V (note: for an electrode, E°ox = E°red)
Cu2+(aq) + 2 e Cu(s); E°red = +0.34V
E°cell = E°red + E°ox. = +0.76 + 0.34 = +1.10V
Electrolysis
Electrolysis involves use of electrical energy to carry out a non-spontaneous redox reaction.
The electrolyte contains ions that are free to move about (and carry the current: in solution,
current involves the movement of charged particles). Cations go to the (-) electrode, the
cathode where reduction takes place. Anions go to the (+) electrode, the anode where
oxidation takes place.
Factors that affect discharge of ions in electrolysis
" Position in the electrochemical series: The lower the metal ion is in the electrochemical
series, the stronger the oxidizing agent, the more readily it gains electrons (is reduced)
at the cathode to form the metal: e.g.: H2 (g) is formed at the cathode instead of Na from
NaCl (aq) or NaOH (aq); however, in electrolysis of CuSO4 (aq), Cu is formed, not H2
at (-) electrode (cathode), since Cu is below H in the activity series.
" Concentration: Sometimes the rule does not match the observed result, e.g., chloride
ions in aqueous solution. For dilute solutions, mainly O2 gas is produced at the (+)
electrode (anode) from the oxidation of water. For concentrated NaCl solution, mainly
Cl2 gas is produced from the oxidation of Cl- at the (+) electrode.
" Nature of the electrode: If an inert electrode is used such as graphite or Pt, it does not
take part in the reactions, e.g., in electrolysis of CuSO4 (aq): H2O is oxidized at the (+)
electrode to produce O2. But if Cu is used as the (+) electrode, the anode, it is oxidized
to Cu2+ ions while Cu2+ ions are reduced to Cu at the (-) electrode and thus [Cu2+]
remains the same.
© IBID Press 2007 2
CHAPTER 9 0XIDATION AND REDUCTION
(TOPICS 9 AND 19) SUMMARY
Electrolysis uses direct current from a battery as a source of electrical energy to carry out a
non-spontaneous redox reaction. The (+) and (-) terminals of the battery are connected to the
electrolytic cell consisting of two electrodes, the (+) electrode, the anode and the (-) electrode,
the cathode, in a solution containing cations (M+) and anions (X-). The battery, a source of
electrons, supplies electrons to the cathode and removes them from the anode. Cations go to
the (-)cathode where reduction takes place to use up the electrons coming in. Anions go to the
(+) anode where oxidation occurs to supply the electrons.
Applications of Electrolysis
Many important metals e.g., Na and Al are prepared by electrolysis of their molten salts; other
chemicals, e.g., NaOH are prepared by electrolysis of aqueous solutions e.g. NaCl (aq).
Electroplating is a process in which a thin layer of metal is deposited on an electrically
conducting surface using electrical energy; it is an important application of electrolytic cells.
In copper plating, for example, reduction occurs at (-) cathode: Cu2+(aq) + 2e- Cu(s);
oxidation occurs at (+) anode: H2O(l) ½O2 (g)+2H+(aq)+2e-. The overall reaction: Cu2+ +
H2O Cu + ½ O2 + 2H+. Thus, as the reaction progresses, pH of solution and color intensity
(due to Cu2+ (aq)) both decrease.
Factors affecting amount of product formed during electrolysis
" Charge on the ion: Na+ (l) + 1 e- Na (l) compared to: Al3+ (l) + 3 e- Al (l): Thus one
mole of Al requires three times the amount of electrons (and thus the charge) compared to 1
mol Na.
" Amount of electrons, i.e., the charge passed through the system: Charge (coulombs) = current
(amperes) x time (seconds). More current used or charge passed for longer time periods
produces more product.
(N.B. Shading indicates Ch 19 (AHL) material.)
© IBID Press 2007 3
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