Electromotive Series of Metals

If from the entire series of standard electrode potentials (see Appendix 11) we separate only the electrode processes that correspond to the general equation:

M^z++ ze^- = M⁰

we get an electromotive (or activity) series of metals. In addition to metals, Hydrogen is included in this series, which permits us to see what metals are capable to displace Hydrogen from aqueous solutions of acids. The electromotive series for the most important metals are given in Table 13.

Table 13. Electromotive Series of Metals

Equation of electrode process	Standard potential at 250C, V	Equation of electrode process	Standard potential at 250C, V
Li+ + → Li0	- 3,045	Co2+ + 2 → Co0	- 0,277
Rb+ + → Rb0	- 2,925	Ni2+ + 2 → Ni0	- 0,250
K+ + → K0	- 2,924	Sn2+ + 2 → Sn0	- 0,136
Cs+ + → Cs0	- 2,923	Pb2+ + 2 → Pb0	- 0,126
Ca2+ + 2 → Ca0	- 2,866	Fe3+ + 3 → Fe0	- 0,037
Na+ + → Na0	- 2,714	2H+ + 2 → H20	0,000
Mg2+ + 2 → Mg0	- 2,363	Bi3+ + 3 → Bi0	0,215
Al3+ + 3 → Al0	- 1,663	Cu2+ + 2 → Cu0	0,337
Ti2+ + 2 → Ti0	- 1,630	Cu+ + → Cu0	0,520
Mn2+ + 2 → Mn0	- 1,179	Hg22+ + 2 → 2 Hg0	0,788
Cr2+ + 2 → Cr0	- 0,913	Ag+ + → Ag0	0,799
Zn2+ + 2 → Zn0	- 0,763	Hg2+ + 2 → Hg0	0,850
Cr3+ + 3 → Cr0	- 0,744	Pt2+ + 2 → Pt0	1,188
Fe2+ + 2 → Fe0	- 0,440	Au3+ + 3 → Au0	1,498
Cd2+ + 2 → Cd0	- 0,403	Au+ + → Au0	1,692

The position of metal in the series characterizes its ability to participate in oxidation-reduction reactions in aqueous solutions in standard conditions. Ions of the metals are oxidizing agents, and the metals in the form of elementary substances are reducing agents. The farther a metal is in the electromotive series, the stronger an oxidizing agent in an aqueous solution is its ions. Conversely, the nearer a metal is to the top of the series the stronger as the reducing properties exhibited by the elementary substance - metal. The potential of the electrode process:

2H⁺ + 2 = H₂

in a neutral medium (pH = 7) is - 0,41 V. The active metals at the top of the series having a potential that is considerably more negative than - 0,41 V displace Hydrogen from water:

2 M + 2H₂O = 2 MOH + H₂↑. Magnesium displaces Hydrogen only from hot water. The metals between Magnesium and Cadmium do not usually displace Hydrogen from water. The surfaces of these metals become covered with oxide films having a protective action.

The metals between Magnesium and Hydrogen displace Hydrogen from solutions of acids. The surfaces of some metals also become covered with protective films that inhibit the reaction. For example, the oxide film on Aluminium makes this metal stable not only in water, but also in solutions of some acids. Lead is not dissolved in Sulfuric acid when its concentration is below 80% because the salt of PbSO₄ is formed when lead reacts with Sulfuric acid is insoluble and produces a protective film on the surface of the metal. The phenomenon of the deep inhibition of the oxidation of the metal due to the presence of protective oxide or salt films on its surface is called passivation, and the state of the metal is called the passive state.

Metals are capable to displace one another from solutions of their salts. The direction of the reaction is determined by their relative position in the electromotive series. When considering specific cases of such reactions, one must remember that active metals displace Hydrogen not only from water, but also from any aqueous solution. Consequently, the mutual displacement of metals from solutions of their salts occurs in practice only with metals below Magnesium in the series. For example:

Zn + CuSO₄ = ZnSO₄ + Cu.

N. Beketov first studied the displacement of metals from their compounds by other metals in detail. As a result of his investigations, he arranged the metals in a "displacement series" according to their chemical activity. This series was the prototype of the electromotive series of metals.

PRACTICE PROBLEMS

1. Determine oxidation number of each element in compounds:

a. CH₄, C₂H₆, C₂H₄, C₂H₂, CH₃OH, C₂H₅OH;

b. CO₂, H₂CO₃, HCOOH, CH₃COOH, H₂C₂O₄;

c. HNO₃, HNO₂, N₂O₄, NO₂, N₂O, Ca₃N₂;

d. NH₃, N₂H₄, NH₂OH, N₂, HNO₂, K₃N;

e. H₂S, H₂SO₃, H₂SO₄, FeS₂, Na₂S₂O₃, Al₂S₃;

f. H₃PO₄, H₃PO₃, PCl₃, P₂O₅, H₄P₂O₇;

g. HClO₄, Ca(ClO)₂, KClO₃, NaClO₂, KCl;

h. CrO, K₃CrO₃, NaCrO₂, BaCrO₄, K₂Cr₂O₇.

2. Balance Redox reactions using method of electron balance. Point out oxidizing and reducing agents:

a. Ca + HNO₃ ® Ca(NO₃)₂ + N₂O + H₂O;

b. K₂S + HNO₃ ® S + NO₂ + KNO₃ + H₂O;

b. NaCrO₂ + Br₂ + NaOH ® Na₂CrO₄ + NaBr + H₂O;

c. HI + KMnO₄ + H₂SO₄ ® I₂ + MnSO₄ + K₂SO₄ + H₂O;

d. Au + HNO₃ + HCl ® AuCl₃ + NO + H₂O;

e. Ag + H₂S + O₂ ® Ag₂S + H₂O.

3. To point out oxidizing and reducing agents and products of reactions. Balance equations:

a. MnO₂ + O₂ + KOH ®

b. Na₂SO₃ + KMnO₄ + H₂SO₄ ®

c. Na₂SO₃ + KMnO₄ + H₂O ®

d. Na₂SO₃ + KMnO₄ + KOH ®

e. NaNO₂ + K₂Cr₂O₇ + H₂SO₄ ®

f. Mg + H₂SO₄ (conc.) ®

g. Fe + HNO₃ (diluted) ®

h. C + HNO₃ (conc.) ®

i. Zn + HNO₃ (diluted) ®

j. P + HNO₃ (diluted) ®

k. HBr + KMnO₄ ®

l. H₂S + HNO₃ (diluted) ®

m. SO₂ + Br₂ + H₂O ®

n. NaNO₂ + K₂Cr₂O₇ + H₂SO₄ ®

o. HBr + HNO₂ ®

p. H₂O₂ + K₂Cr₂O₇ + H₂SO₄ ®

q. SO₂ + Br₂ + H₂O ®

4. Point out a type of redox reaction (Intermolecular, Intramolecular, Disproportionation). Determine chemical elements charged oxidation degree. To balance equations:

a. Ca(OH)₂ + Cl₂ ® CaCl₂ + Ca(ClO)₂ + H₂O;

b. NO₂ + H₂O ® HNO₃ + NO;

c. KClO ® KClO₃ + KCl;

d. H₃PO₃ ® PH₃ + H₃PO₄;

e. HCl + KClO₃ ® Cl₂ + KCl + H₂O;

f. H₂S + H₂SO₄ ® SO₂ + H₂O;

g. KI + KIO₃ + H₂SO₄ ® I₂ + K₂SO₄ + H₂O;

h. KClO₃ + H₂SO₄ ® ClO₂ + O₂ + K₂SO₄ + H₂O;

i. HNO₃ ® NO₂ + O₂ + H₂O;

j. Pb(NO₃)₂ ® PbO + NO₂ + O₂;

k. (NH₄)₂Cr₂O₇ ® N₂ + Cr₂O₃ + H₂O;

l. Au(NO₃)₃ ® Au + NO₂ + O₂;

m. H₂O₂ ® H₂O + O₂.

Date: 2015-01-12; view: 1978