Chemical Forums
Chemistry Forums for Students => Physical Chemistry Forum => Topic started by: mistertaylor92 on March 02, 2012, 06:52:19 PM
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For a lab we constructed an electrochemical cell that consists of two compartments; one consisting of a zinc electrode submersed in a 0.15 M zinc chloride solution (the anode) and another consisting of a platinum electrode submersed in a solution that is 0.5 M in both potassium ferrocyanide (K4Fe(CN)6•3H2O) and potassium ferricyanide (K3Fe(CN)6) (the cathode). I need to calculate Q for the reaction.
From what I understand, the representation of the cell is as follows:
Zn(s) | Zn2+aq(0.15M) || Fe(CN)63-(aq)(0.10 M), Fe(CN)64-(aq)(0.10 M) | Pt
And the overall reaction is as follows:
Zn(s) + 2Fe(CN)64-(aq) :rarrow: Zn2+aq + 2Fe(CN)63-(aq)
Am I right to say that the reaction quotient for this reaction is equal to:
Q = [Zn2+]*[Fe(CN)63-]2/[Fe(CN)64-]2
= (0.15 M)(0.10 M)2/(0.10 M)2
= 0.15
The lab manual states that activities can be ignored. Thanks in advance.
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In general looks good, but why 0.1 if 0.5?
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My apologies that was a typo. The molarity of the iron compounds are both 0.1 M.