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If temperature of the cell $ Zn | Zn^{2+}_{C_1} || Sn^{2+}_{C_2} | Sn (C_1 < C_2) $ decreases then theoritically cell potential what changes are observed in the cell potential of the cell?
What is the equilibrium constant of the reaction taking place in the cell $ Zn | Zn^{2+}_{(C-1)} || Cu^{2+}_{(C-2)} | Cu at 25 ^\circ C? ( E ^ \circ _{Zn ^{2+} |Zn} = − 0.76 V., E ^\circ Cu ^{ 2+} | Cu= 0.34 V.)$
The equilibrium constant (K) of an electrochemical cell is related to the standard electrode potential (E°) by the equation: K = exp(nFE°/RT), where n is the number of electrons transferred, F is the Faraday constant, R is the gas constant, and T is the absolute temperature. Given E°(Zn²âº/Zn) = -0.76 V and E°(Cu²âº/Cu) = 0.34 V, we can calculate E° for the overall cell reaction: E°cell = E°cathode - E°anode = 0.34 V - (-0.76 V) = 1.10 V. Substituting the values, we get K = exp(2 × 96485 C/mol × 1.10 V / (8.314 J/mol·K × 298 K)) = 1.9413 × 10³â·.
If for the reaction taking place in one electrochemical cell $ Kc = 1.3141 \times 10^{107} M^{– 1} $ and n = 2 then for that cell value of E ^\circ $ Cell is _.
If for the reaction taking place in an electrochemical cell $ c = 1.3141 \times 10^{307} $ and n = 6. if standard reduction potential of the cathode is 1.36 V then what is the standard oxidation potential of the anode?
What is the pH of the solution of HCl if cell potentila of the cell $ Pt | H_{2(g. 1.0 bar)}| HCl_{(xM)} || Au ^{3+}_{(0.001M)} | Au is 1.6655 V ? (E^0_{Au^{3+}|Au} = 1.4 V.) $
Mention the value of n for the reaction taking place in the cell if its cell potential is 3.16V and equilibrium constant is $ 1.727 \times 10^{214}M ^{-2} at 25 ^ \circ C $ temperature.
If for the cell $ Mg | Mg^{2+}_{(0.2M)} || H^+_{(xM)}∣ H_{2(g,1.0 bar )}|Pt $ values of E and $ E^0_{cell} are 2.3629 V.(0.2 M)(X M)Cell and 2.36 V. respectively at 25 ^\circ C $ temperature then value of x is _.
If standard potential of the reaction $ 2H_2O_(l) \rightleftharpoons O_{2(g,1 bar )} + 4H^+_{aq} + 4e^− is -1.23V and [H+] = 10^{-7} at 25 ^\circ C $ temperature then what is the value of potential for the reaction? OR What is the oxidation potential of the pure water?
The standard potential of the given reaction is -1.23V. At [H+] = 10^-7 M, the potential is given by the Nernst equation: E = E° - (0.0592/4) log (10^-28). Substituting the values, E = -1.23 - (0.0592/4)(-28) = -0.817V.
If for the reaction $ 2H_2O_(l) + 2 e ^ - \rightleftharpoons H_{2 (g, 1 bar)} + 2 OH ^ - _{aq, 10 ^ {-7} M } E = -0.417 V at 25 ^ \circ C $ temperature what is the value of E 0 for the reaction? OR what is the standard reduction potential of the pure water at $ 25 ^\circ C $ if potential is – 0.417 V.
The given potential of -0.417V corresponds to [H+] = 10^-7 M. To find E°, we use the Nernst equation: E° = E + (0.0592/2) log [H+]. Substituting the values, E° = -0.417 + (0.0592/2)(-7) = -0.83V, which is the standard reduction potential of water.
If $ Co | Co^{3+}, Co^{2+}| Co ^{3+}, Fe | Fe ^ {3+}, Fe^{2+}| Fe^{3+} std. oxidation potential are − 0.4167 V., − 1.81 V., 0.0367 V., − 0.77 V. respectively then what is the std. potential of the cell Fe | Fe^{2+} || Co^{2+} Co $ ?
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