Question

Given below are the half-cell reactions

Mn²⁺ +2e^- →Mn, E^o= -1.18V

2(Mn³⁺ +e^- →Mn²⁺), E^o=+1.51V

The E⁰ for 3Mn²⁺ → Mn +2Mn³⁺ will be

(a) -2.69V the reaction will not occur

(b) -2.69 V the reaction will occur

(c) -0.33V, the reaction will not occur

(d) -0.33V, the reaction will occur

Answer 1

Correct option (a) -2.69V the reaction will not occur

Explanation:

Alternate method

Comments

I have doubt over this .In the second  equation we are multiplying with 2 right in order to get equation 3 .then why we dont multiplythe value of E0 of second  equation. Plz answer it.

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Because Eo is standard electrode potential that measures the individual potential of the reversible electrode at standard state with ions at an effective concentration and due to electrolytic cell behaviour we can't multiply it by E0

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I have a doubt that, why
the E0 = -1.18-1.51  as in many questions  
the greater value of E0 is consider as cathode and smaller value as anode . Then why here we have taken greater as anode and smaller as cathode??

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that's why the reaction is non-spontaneous...