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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:

Mz++ 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: 1127


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