Introduction: Many elements possess more than one oxidation state. These include representatives of both the main group elements (the s- and p-blocks, such as N, S, and P) and the transition elements (the d- and f- blocks, such as Fe, Co, and Cr). Manganese, for example, is known in the 0, II, III, IV, V, VI, and VII oxidation states. In most cases for transition metals, each oxidation state exhibits a characteristic color. In this experiment, the more accessible oxidation states of manganese will be investigated by reducing the Mn(VII) in permanganate ion (MnO4-) in a series of redox titrations. In a reduction, the material being reduced gains electrons, and consequently has its oxidation state lowered. In acid solution, for example, Mn(VII) is reduced to Mn(I) (manganous ion) through the gain of five electrons, as shown in the following unbalanced equation: Mno, + 5 e- + H* 2 Mn²+ These equations are called half-reactions because only one species is changing oxidation state, requiring electrons in the balanced equation. In neutral solution, Mn(VII) is reduced to MnO2 (Mn(IV), manganese dioxide), and in basic solution, Mn(VII) is reduced to Mn(VI) (MnO,2-, manganate ion): Mno, + 3 e¯ 2 MnO2 MnO, + e- + OH¯2 Mn0?- Since the electrons gained by the Mn(VII) must come from somewhere, for every material being reduced, some species must simultaneously be oxidized. In an oxidation, the material being oxidized loses electrons, and consequently has its oxidation state increased. In this experiment, for example, it is the bisulfite ion (HSO3-) that is oxidized to sulfate ion (SO4²-): HSO3 2 S03- + 2 e-
Introduction: Many elements possess more than one oxidation state. These include representatives of both the main group elements (the s- and p-blocks, such as N, S, and P) and the transition elements (the d- and f- blocks, such as Fe, Co, and Cr). Manganese, for example, is known in the 0, II, III, IV, V, VI, and VII oxidation states. In most cases for transition metals, each oxidation state exhibits a characteristic color. In this experiment, the more accessible oxidation states of manganese will be investigated by reducing the Mn(VII) in permanganate ion (MnO4-) in a series of redox titrations. In a reduction, the material being reduced gains electrons, and consequently has its oxidation state lowered. In acid solution, for example, Mn(VII) is reduced to Mn(I) (manganous ion) through the gain of five electrons, as shown in the following unbalanced equation: Mno, + 5 e- + H* 2 Mn²+ These equations are called half-reactions because only one species is changing oxidation state, requiring electrons in the balanced equation. In neutral solution, Mn(VII) is reduced to MnO2 (Mn(IV), manganese dioxide), and in basic solution, Mn(VII) is reduced to Mn(VI) (MnO,2-, manganate ion): Mno, + 3 e¯ 2 MnO2 MnO, + e- + OH¯2 Mn0?- Since the electrons gained by the Mn(VII) must come from somewhere, for every material being reduced, some species must simultaneously be oxidized. In an oxidation, the material being oxidized loses electrons, and consequently has its oxidation state increased. In this experiment, for example, it is the bisulfite ion (HSO3-) that is oxidized to sulfate ion (SO4²-): HSO3 2 S03- + 2 e-
Introduction: Many elements possess more than one oxidation state. These include representatives of both the main group elements (the s- and p-blocks, such as N, S, and P) and the transition elements (the d- and f- blocks, such as Fe, Co, and Cr). Manganese, for example, is known in the 0, II, III, IV, V, VI, and VII oxidation states. In most cases for transition metals, each oxidation state exhibits a characteristic color. In this experiment, the more accessible oxidation states of manganese will be investigated by reducing the Mn(VII) in permanganate ion (MnO4-) in a series of redox titrations. In a reduction, the material being reduced gains electrons, and consequently has its oxidation state lowered. In acid solution, for example, Mn(VII) is reduced to Mn(I) (manganous ion) through the gain of five electrons, as shown in the following unbalanced equation: Mno, + 5 e- + H* 2 Mn²+ These equations are called half-reactions because only one species is changing oxidation state, requiring electrons in the balanced equation. In neutral solution, Mn(VII) is reduced to MnO2 (Mn(IV), manganese dioxide), and in basic solution, Mn(VII) is reduced to Mn(VI) (MnO,2-, manganate ion): Mno, + 3 e¯ 2 MnO2 MnO, + e- + OH¯2 Mn0?- Since the electrons gained by the Mn(VII) must come from somewhere, for every material being reduced, some species must simultaneously be oxidized. In an oxidation, the material being oxidized loses electrons, and consequently has its oxidation state increased. In this experiment, for example, it is the bisulfite ion (HSO3-) that is oxidized to sulfate ion (SO4²-): HSO3 2 S03- + 2 e-
Balance the three redox reactions in this experiment by combining each of the three permanganate half-reactions on page 1 of the Introduction with the bisulfite half-reaction. One should be in acid, one should be in neutral solution, and another should be in base as indicated.
Transcribed Image Text:Introduction:
Many elements possess more than one oxidation state. These include representatives of both the main
group elements (the s- and p-blocks, such as N, S, and P) and the transition elements (the d- and f-
blocks, such as Fe, Co, and Cr). Manganese, for example, is known in the 0, II, III, IV, V, VI, and VII
oxidation states. In most cases for transition metals, each oxidation state exhibits a characteristic color.
In this experiment, the more accessible oxidation states of manganese will be investigated by reducing
the Mn(VII) in permanganate ion (MnO4-) in a series of redox titrations. In a reduction, the material
being reduced gains electrons, and consequently has its oxidation state lowered. In acid solution, for
example, Mn(VII) is reduced to Mn(I) (manganous ion) through the gain of five electrons, as shown in
the following unbalanced equation:
Mno, + 5 e- + H* 2 Mn²+
These equations are called half-reactions because only one species is changing oxidation state, requiring
electrons in the balanced equation. In neutral solution, Mn(VII) is reduced to MnO2 (Mn(IV), manganese
dioxide), and in basic solution, Mn(VII) is reduced to Mn(VI) (MnO,2-, manganate ion):
Mno, + 3 e¯ 2 MnO2
MnO, + e- + OH¯2 Mn0?-
Since the electrons gained by the Mn(VII) must come from somewhere, for every material being
reduced, some species must simultaneously be oxidized. In an oxidation, the material being oxidized
loses electrons, and consequently has its oxidation state increased. In this experiment, for example, it is
the bisulfite ion (HSO3-) that is oxidized to sulfate ion (SO4²-):
HSO3 2 S03- + 2 e-
Method Method Redox reactions are generally balanced using two processes. The first method is based on the redox reaction’s division into two reactions where one is involved in oxidation and the other in reduction (half reaction method) and the second one is based on the reducing and oxidizing molecule’s oxidation number variation (oxidation number method).
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