The standard electrode potentials at 298 K
Ions are present as aqueous species and H ₂O as liquid; gases and solids are shown by g and s respectively.

	Reaction (Oxidised form + ne- \(\rightarrow\) Reduced form)		\(\mathbf{E}^{\odot} / \mathbf{V}\)
	\(\mathbf{F}_{2}(g)+2 e^{-} \longrightarrow 2 \mathbf{F}\) \(\mathrm{Co}^{3+}+e^{-} \longrightarrow \mathrm{Co}^{2+}\) \(\mathbf{H}_{2} \mathbf{O}_{2}+\mathbf{2} \mathbf{H}^{+}+\mathbf{2} e^{-} \longrightarrow \mathbf{2} \mathbf{H}_{2} \mathbf{O}\) \(\mathrm{MnO}_{4}^{-}+8 \mathrm{H}^{+}+5 e^{-} \longrightarrow \mathrm{Mn}^{2+}+4 \mathrm{H}_{2} \mathrm{O}\) \(\mathbf{A u}^{3+}+\mathbf{3} \boldsymbol{e}^{-} \longrightarrow \mathbf{A u}(s)\) \(\mathrm{Cl}_{2}(g)+2 e^{-} \longrightarrow 2 \mathrm{Cl}^{-}\) \(\mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}+\mathbf{1 4 H}^{+}+6 e^{-} \longrightarrow 2 \mathrm{Cr}^{3+}+7 \mathrm{H}_{2} \mathrm{O}\) \(\mathrm{O}_{2}(\mathrm{~g})+4 \mathrm{H}^{+}+4 e^{-} \longrightarrow 2 \mathrm{H}_{2} \mathrm{O}\) \(\mathrm{MnO}_{2}(s)+4 \mathrm{H}^{+}+2 e^{-} \longrightarrow \mathrm{Mn}^{2+}+2 \mathrm{H}_{2} \mathrm{O}\) \(\mathbf{B r}_{2}+2 e^{-} \longrightarrow 2 \mathbf{B r}^{-}\) \(\mathrm{NO}_{3}^{-}+4 \mathrm{H}^{+}+3 e^{-} \longrightarrow \mathrm{NO}+\underline{2} \mathrm{H}_{2} \mathrm{O}\) \(2 \mathrm{Hg}^{2+}+2 e^{-} \longrightarrow \mathrm{Hg}_{2}^{2+}\) \(\mathbf{A g}^{+}+\boldsymbol{e}^{-} \longrightarrow \mathbf{A g}(\boldsymbol{s})\) \(\mathbf{F e}^{3+}+e^{-} \longrightarrow \mathbf{F e}^{2+}\) \(\mathrm{O}_{2}(g)+2 \mathrm{H}^{+}+2 e^{-} \longrightarrow \mathbf{H}_{2} \mathrm{O}_{2}\) \(\mathbf{I}_{2}(s)+2 e^{-} \longrightarrow 2 \mathbf{I}^{-}\) \(\mathbf{C u}^{+}+e^{-} \longrightarrow \mathbf{C u}(\boldsymbol{s})\) \(\mathrm{Cu}^{2+}+2 e=\mathrm{Cu}(s)\) \(\mathbf{A g C l}(s)+e^{-} \longrightarrow \mathbf{A g}(s)+\mathbf{C l}\) \(\mathbf{A g B r}(s)+e^{-} \longrightarrow \mathbf{A g}(s)+\mathbf{B r}^{-}\) \(2 \mathrm{H}^{+}+2 e^{-} \longrightarrow \mathbf{H}_{2}(g)\) \(\mathbf{P b}^{2+}+2 e^{-} \longrightarrow \mathbf{P b}(s)\) \(\mathrm{Sn}^{2+}+2 e^{-} \longrightarrow \operatorname{Sn}(s)\)\(\)\(\) \(\)\(\mathbf{N i}^{2+}+2 e^{-} \longrightarrow \mathbf{N i}(s)\) \(\)\(\mathrm{Fe}^{2+}+2 e-\mathrm{Fe}(s)\) \(\)\(\mathrm{Cr}^{3+}+3 e^{-} \longrightarrow \mathrm{Cr}(s)\) \(\)\(\mathbf{Z n}^{2+}+2 e^{-} \longrightarrow \mathbf{Z n}(s)\) \(\)\(\mathbf{2} \mathbf{H}_{2} \mathbf{O}+\mathbf{2} e^{-} \longrightarrow \mathbf{H}_{2}+\mathbf{2} \mathbf{O H}^{-}\) \(\)\(\mathbf{A l}^{3+}+\mathbf{3 e}^{-} \longrightarrow \mathbf{A l}(\boldsymbol{s})\) \(\)\(\mathbf{M g}^{2+}+2 e^{-} \longrightarrow \mathbf{M g}(s)\) \(\)\(\)\(\mathrm{Na}^{+}+e^{-} \longrightarrow \mathrm{Na}(s)\) \(\mathrm{Ca}^{2+}+2 e^{-} \longrightarrow \mathrm{Ca}(s)\) \(\mathbf{K}^{+}+e^{-} \longrightarrow \mathbf{K}(\boldsymbol{s})\) \(\mathbf{L i}^{+}+e^{-} \longrightarrow \mathbf{L i}(s)\)		2.87
			1.81
			1.78
			1.51
			1.40
			1.36
			1.33
			1.23
			1.23
			1.09
			0.97
			0.92
			0.80
			0.77
			0.68
			0.54
			0.52
			0.34
			0.22
			0.10
			0.00
			0.13
			0.14
			0.25
			0.44
			0.74
			0.76
			0.83
			1.66
			2.36
			2.71
			2.87
			2.93
			3.05

1. A negative \(\boldsymbol{E}^{\ominus}\) means that the redox couple is a stronger reducing agent than the H⁺/H₂ couple.
2. A positive \(\boldsymbol{E}^{\ominus} \) means that the redox couple is a weaker reducing agent than the H⁺/H₂ couple.
(a) Which is best oxidising agent and why?
(b) Which is best reducing agent and why?
(c) Calculate \(\mathbf{E}^{\circ} \text { cell }\) lof Zn(s)/Zn²⁺(1M) II Cu^{2 +}(aq)/Cu(s).
(d) If a redox couple has negative E^o, is it stronger or weaker reducing agent than H⁺/H₂ couple?
(e) Arrange K, Ag, Hg, Mg, Cr in increasing order of reducing power.
(f) Arrange AI, Cu, Fe, Mg and Zn in order in which they will displace each other from salt solution.

Redox reactions are important class of reactions which are taking place in our daily life. Metals are good reducing agents because they can lose electrons easily whereas non-metals are good oxidising agents which can gain electrons easily. In electrolytic cells, electricity is passed to bring about redox reaction. All rechargeable batteries act as electrolytic cells while recharging. Electrochemical cells produce electricity as a result of redox reaction. Salt bridge is used in electrochemical cell to complete internal circuit and prevents accumulation of charges.
(a) What is electrochemical cell?
(b) Why is anode called oxidation electrode?
(c) Give one example of rechargeable cells widely used in vehicles.
(d) Highly reactive metals are obtained by electrolysis of their molten ores, why?
(e) What is direction of flow of current and electrons?
(f) What is standard electrode potential?
(g) What is meaning of -ve value of reduction potential? Give example.

Redox reactions are reactions in which oxidation and reduction takes place simultaneously. Oxidation number are assigned in accordance with the set of rules. Oxidation number and ion electron methods both are used in balancing ionic equations. Redox reactions are classified as combination, decomposition, displacement and disproportionation reactions. The concept of redox couple and electrode processes is basis of electrolysis and electrochemical cells.
(a) What are oxidation number of each individual Br in Br₃ O₈ ?
(b) If electrolysis of CuSO₄ solution is carried out using Cu electrodes, what will be reaction taking place at anode.
(c) What is oxidation number of Cr in CrO₅?
(d) Give one example of disproportionation reaction.