If I compare Cu and Zn from the periodic table,
Cu--3d10, 4s1 -- means can easily loose e, thus should undergo oxidation easily.
Zn --3d10, 4s2---full valence shell, therefore loosing e is not easy, therefore shouldnot undergo oxidation.
But in real electrochem cells Zn undergoes oxidation instead of Cu. Why?
Also, I compared Cu and Ag--4d10 5s1--Ag being bigger in size should loose electons easily as compared to Cu and should undergo oxidation easily, but this is not observed. Infact, Ag undergoes reduction in presence of Cu.
I am also aware of the table of strong reducing agents. Zn is placed higher than cr, Ni,Cu, Ag. But I fail to understand how is this table arranged? I tried to find a "link" between this table and the periodic table or the distribution of the valence e. It is not making much sense to me.
Li, K, Ca etc are on the left of periodic table and are good reducing agents=undergo oxidation=loose e.
Also Cu is on the left of Zn then why it doesn't undergo oxd?
2007-04-27
06:15:25
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1 answers
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AS
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Chemistry