I'm working on a problem right now: What is the equilibrium constant for the oxidation of copper by H+ for the following reaction: Cu(s) + 2H+(aq) >>> Cu2+(aq) + H2(g). I know the E cell value for the reaction the way it is set up is -.34, so I plug that value into the Nernst equation for the standard value, but why does the E cell value on the other side of the Nernst equation equal zero?
Ecell = Ecell(standard e cell) -((RT)/(vEF))ln(Q)
What I'm trying to find is Q. But what I don't understand is why the Ecell on left side of the equation as it's written is zero. I know the E(v) for hydrogen is zero, but is this why the Ecell value is zero? Because energy must be supplied at conditions that aren't standard? Thanks for the help.
Ecell = Ecell(standard e cell) -((RT)/(vEF))ln(Q)
What I'm trying to find is Q. But what I don't understand is why the Ecell on left side of the equation as it's written is zero. I know the E(v) for hydrogen is zero, but is this why the Ecell value is zero? Because energy must be supplied at conditions that aren't standard? Thanks for the help.
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I have a little trouble following the question, so I'll just make some general comments.
E zero is the voltage when everything is at standard state.
You can calculate the actual E at any set of concentrations by plugging those values into Q.
If the actual E is zero, then you have equilibrium conditions and, instead of Q, you have Keq
E zero is the voltage when everything is at standard state.
You can calculate the actual E at any set of concentrations by plugging those values into Q.
If the actual E is zero, then you have equilibrium conditions and, instead of Q, you have Keq