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Oxidation and Reduction
Oxidation number:
Fictive charges which are given to an element by some rules. The oxidation number is
described by positive or negative roman numbers and they are written above or in brackets
after the element they are assigned to.
Rules of the assignment of the oxidation number:
*the sum of the ON’s of all the atoms in a compound is equal to its total charge
*atoms in their elemental form have ON = ±0: Na, Cl2, Mn,…
*atoms of group 1+2 have in compounds the number of their group as ON: +I/+II
*hydrogen: +I in compounds with nonmetals, -I in compounds with metals
*fluorine: -I in all its compounds (because of highest electronegativity (4.0))
*oxygen: -II (except in compounds with F), -I in peroxides (O22-), -1/2 in superoxides (O2-), 1/3 in ozonides (O3-)
*halogens: -I in most compounds, unless the other elements include oxygen ore more
electronegative halogens
*in compounds which are built out of non-metals, the oxidation number of the more
electronegative element is equal to the charge of its anion. The other part of the compound has
a positive oxidation number Æsum of the ON in an uncharged compound has to be 0
Determination of the oxidation number of a compound:
*draw the Lewis-structure and split the binding electrons heterolytically (by use of
electronegativity)
* determine the oxidation number via a comparison of valence electrons of the considered
atom with the number of electrons after heterolytic splitting
*is the number of determined electrons bigger than the number of valence electrons of the
element, the atom has a negative oxidation number (less electrons Æ positive ox.-number)
Definitions of oxidation and reduction:
*old definition: an element/ compound is oxidized, when it is treated with oxygen. An
element/ compound is reduced when it loses oxygen.
*new definition: an element/ compound is oxidized when it gets a higher/ more positive
oxidation number during a reaction, that means it loses electrons (unequal to reduction)
Redox-reactions:
*a chemical reaction is a redox-reaction, when at least some of the oxidation numbers of the
elements/ compounds change Æ there is always also an oxidation when there is a reduction
*to solve the equation of a redox-reaction, determine the oxidation numbers of each element,
define possible products, divide the reaction into reduction and oxidation, count the initial and
the final ON of each element and add missing e-, equal the numbers of electrons of the
reduction and oxidation, combine both equations to the final redox-equation
Example 1
xKMnO4 + yH2O2 + zH2SO4 Æ ?
Oxidation numbers: Mn(+VII), O(-II in H2SO4 and -I in H2O2), K(+I), H(+I), S(+VI);
Mn(+VII) is reduced to Mn(+II) in an acidic medium; O can be oxidized (-I to ±0);
Æ possible products: Mn2+, O2
xKMnO4 + yH2O2 + zH2SO4 Æ aMnSO4 + bK2SO4 + cH2O + dO2
Reduction:
Oxidation:
Redox:
Mn(+VII) + 5e- Æ Mn(+II)
Æ O2 (each ±0) + 2eO22-(each –I)
22Mn(+VII) + 5O2
Æ Mn(+II) + 5O2
×2
×5
Æ 2KMnO4 + 5H2O2 + 3H2SO4 ô 2MnSO4 + K2SO4 + 8H2O + 5O2
Example 2
xCr2O72- + ySO32- + zH3O+ Æ ?
Oxidation numbers: Cr(+VI), O(-II), S(+IV), H(+I);
Cr could be reduced from +VI to +III, S is oxidized from +IV to +VI;
Æ possible products: Cr3+, SO42xCr2O72- + ySO32- + zH3O+ Æ aCr3+ + bSO42- + cH2O
Reduction:
Oxidation:
Redox:
2Cr(+VI) + 6e- Æ 2Cr3+
×1
22Æ SO4 + 2e
×3
SO3
2Cr(+VI) + 3SO32- Æ 2Cr3+ + 3SO42-
Æ Cr2O72- + 3SO32- + 8H3O+ ô 2Cr3+ + 3SO42- + 12H2O
Example 3
xBrO3- + yFe2+ + zH3O+ Æ ?
Oxidation numbers: Br(+V), Fe2+(+II), O(-II), H(+I);
Br(+V) could be reduced e.g. to Br-, Fe2+ could be oxidized to Fe3+,
Æ possible products: Fe3+, BrxBrO3- + yFe2+ + zH3O+ Æ aBr- + bFe3+ + cH2O
×1
Reduction:
BrO3- + 6e- Æ Br- + 3O22+
3+
Oxidation:
Fe Æ Fe + 1e
×6
Redox: BrO3- + 6Fe2+ Æ Br- + 6Fe3+ + 3O2Æ BrO3- + 6Fe2+ + 6 H3O+ ô Br- + 6Fe3+ + 9H2O
Example 4
xCa3(PO4)2+yC+zSiO2 Æ ?
Oxidation numbers: Ca2+(+II), P(+V), O(-II), C(±0), Si(+IV)
P could be reduced, e.g. to ±0, C could be oxidized, e.g. to +II
Æ possible products: P, CO
xCa3(PO4)2 + yC + zSiO2 Æ aCaSiO3 + bCO + cP
Reduction:
2PO43- + 10e- Æ 2P + 8O2Oxidation:
C Æ CO + 2eRedox:
2PO43- + 5C Æ 5CO + 2P + 8O2-
×1
×5
Æ Ca3(PO4)2 + 5C + 3SiO2 ô 3CaSiO3 + 5CO + 2P
Literature:
Shriver/Atkins, Inorganic Chemistry, Fourth Edition, Oxford University Press, 2006, 39-40
Shriver/Atkins, Inorganic Chemistry, Fourth Edition, Oxford University Press, 2006, 141ff