Electrochmistry

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NOTES:-

*for F2 gas , SEP is the highest in the table indicating that F2 has the highest tendency to get reduced to F- : F2 is the strongest oxidising agent.

*Li has the lowest SEP indicating that Li+ is the weakest oxidising agent and Li metal is the strongest reducing agent.

*As value of SRP decreases for metal ion, reducing power metal increases.

Applications of Electrochemical Series

(i) Reactivity of metals: The activity of the metal depends on its tendency to lose electron or electrons, i.e., tendency to form cation . This tendency depends on the magnitude of standard reduction potential. The metal which has high negative value (or smaller positive value) of standard reduction potential readily loses the electron or electrons and is converted into cation. Such a metal is said to be chemically active. The chemical reactivity of metals decreases from top to bottom in the series. The metal higher in the series is more active than the metal lower in the series. For example,

(a) Alkali metals and alkaline earth metals having high negative values of standard reduction potentials are chemically active. These react with cold water and evolve hydrogen. These readily dissolve in acids forming corresponding salts and combine with those substances which accept electrons.

(b) Metals like Fe, Pb, Sn, Ni, Co, etc., which lie a little down in the series do not react with cold water but react with steam to evolve hydrogen.

(c) Metals like Cu, Ag and Au which lie below hydrogen are less reactive and do not evolve hydrogen from water.

(ii) Electropositive character of metals :The electropositive character also depends on the tendency to lose electron or electrons. Like reactivity, the electropositive character of metals decreases from top to bottom in the electrochemical series. On the basis of standard reduction potential values, metals are divided into three groups

(a) Strongly electropositive metals : Metals having standard reduction potential near about – 2.0 volt or more negative like alkali metals, alkaline earth metals are strongly electropositive in nature.

(b) Moderately electropositive metals : Metals having values of reduction potentials between 0.0 and about – 2.0 volt are moderately electropositive Al, Zn, Fe, Ni, Co, etc., belong to this group.

(c) Weakly electropositive : The metals which are below hydrogen and possess positive values of reduction potentials are weakly electropositive metals. Cu, Hg, Ag, etc., belong to this group.

(iii) Displacement reactions

(a) To predict whether a given metal will displace another, from its salt solution: A metal higher in the series will displace the metal from its solution which is lower in the series, i.e., The metal having low standard reduction potential will displace the metal from its salt solution which has higher value of standard reduction potential. A metal higher in the series has greater tendency to provide electrons to the cations of the metal to be precipitated.

(b) Displacement of one non metal from its salt solution by another non metal: A non-metal higher in the series (towards bottom side), i.e., having high value of reduction potential will displace another non-metal with lower reduction potential, i.e., occupying position above in the series. The non-metals which possess high positive reduction potentials have the tendency to accept electrons readily. These electrons are provided by the ions of the nonmetal having low value of reduction potential,. Thus, can displace bromine and iodine from bromides and iodides.

[The activity or electronegative character or oxidising nature of the nonmetal increases as the value of reduction potential increases.]

(c) Displacement of hydrogen from dilute acids by metals : The metal which can provide electrons to ions present in dilute acids for reduction, evolve hydrogen from dilute acids.The metal having negative values of reduction potential possess the property of losing electron or electrons.Thus, the metals occupying top positions in the electrochemical series readily liberate hydrogen from dilute acids and on descending in the series tendency to liberate hydrogen gas from dilute acids decreases.The metals which are below hydrogen in electrochemical series like Cu, Hg, Au, Pt, etc., do not evolve hydrogen from dilute acids.

(d) Displacement of hydrogen from water : Iron and the metals above iron are capable of liberating hydrogen from water. The tendency decreases from top to bottom in electrochemical series. Alkali and alkaline earth metals liberate hydrogen from cold water but Mg, Zn and Fe liberate hydrogen from hot water or steam.

(iv) Reducing power of metals: Reducing nature depends on the tendency of losing electron or electrons. More the negative reduction potential, more is the tendency to lose electron or electrons. Thus reducing nature decreases from top to bottom in the electrochemical series. The power of the reducing agent increases, as the standard reduction potential becomes more and more negative. Sodium is a stronger reducing agent than zinc and zinc is a stronger reducing agent than iron. (decreasing order of reducing nature.Alkali and alkaline earth metals are strong reducing agents.

(v) Oxidising nature of non-metals : Oxidising nature depends on the tendency to accept electron or electrons. More the value of reduction potential, higher is the tendency to accept electron or electrons. Thus, oxidising nature increases from top to bottom in the electrochemical series. The strength of an oxidising agent increases as the value of reduction potential becomes more and more positive.

(vi) Thermal stability of metallic oxides : The thermal stability of the metal oxide depends on its electropositive nature. As the electropositivity decreases from top to bottom, the thermal stability of the oxide also decreases from top to bottom. The oxides of metals having high positive reduction potentials are not stable towards heat. The metals which come below copper form unstable oxides, i.e., these are decomposed on heating.

(vii) Extraction of metals : A more electropositive metal can displace a less electropositive metal from its salt solution.

Determining Equilibrium Constant with Nernst Equation

When the reactants and the products of the electrochemical cell reach equilibrium, the value of ΔG becomes 0. At this point, the reaction quotient and the equilibrium constant (Kc) are the same. Since ΔG = -nFE, the cell potential at equilibrium is also 0.

Substituting the values of Q and E into the Nernst equation, the following equation is obtained.

0 = E0cell – (RT/nF) ln Kc

The relationship between the Nernst equation, the equilibrium constant, and Gibbs energy change is illustrated below.

Nernst equation vs Equilibrium constant vs Gibbs energy change

Converting the natural logarithm into base-10 logarithm and substituting T=298K (standard temperature), the equation is transformed as follows.

E0cell = (0.0592V/n) log Kc

Rearranging this equation, the following equation can be obtained.

log Kc = (nE0cell)/0.0592V

Thus, the relationship between the standard cell potential and the equilibrium constant is obtained. When Kc is greater than 1, the value of E0cell will be greater than 0, implying that the equilibrium favours the forward reaction. Similarly, when Kc is less than 1, E0cell will hold a negative value which suggests that the reverse reaction will be favoured.

Variation of Molar Conductivity:-

i) Nature of electrolytes:-

a) Strong electrolytes have higher values of Molar Conductivity.[Eg: HCl, NaCl, NaOH]

b) Weak electrolytes have lower values of Molar Conductivity.[ Eg: CH3COOH, NH4OH, etc]

ii) Concentration of Solution:-

The Molar Conductivity increases with dilution.

iii) Temperature:-

As temperature increases, the molar conductivity also increases.

Kohlrausch's law

Some important applications of Kohlrausch law of independent migration of ions are:

1.Kohlrausch law helps us in the determination of limiting molar conductivities for any electrolyte. Weak electrolytes have lower molar conductivities and lower degree of dissociation at higher concentrations. The graph plotted between molar conductivity and c1/2 (where c is the concentration) is not a straight line for weak electrolytes. The molar conductivity of weak electrolyte increases steeply at lower concentrations. Therefore, limiting molar conductivity, Λ cannot be obtained by extrapolation of molar conductivity to zero concentration. Hence, we use the Kohlrausch law of independent migration of ions for the determination of limiting molar conductivity, Λ for weak electrolytes.

2.Kohlrausch law also helps us in determining the value of dissociation constant from the value of molar conductivity and limiting molar conductivity for a weak electrolyte at a given concentration.

α = ΛËm∘

Where, α = dissociation constant

Λ = molar conductivity

Ëm° = limiting molar conductivity