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Question

The reaction of potassium permanganate with acidified iron [II] sulphate is given below :- 2KMnO4+10FeSO4+8H2SO4K2SO4+2MnSO4+5Fe2(SO4)3+8H2O
If 15.8 g of potassium permanganate was used in the reaction, calculate the mass of iron [II] sulphate used in the above reaction.
[Molecular Weight of K=39,Mn=55,Fe=56,S=32,O=16 g/mole]

A
76 g
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B
152 g
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C
38 g
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D
46 g
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Solution

The correct option is A 76 g

Molecular weight of KMnO4=39+55+164=158

Molecular weight of K2SO4=239+32+164=174

Molecular weight of FeSO4=56+32+164=152

2×158 g of KMnO4 yields 174 g of K2SO4

15.8 g of KMnO4 will yield [(174)/(2×158)]×15.8 g = 8.7 g of K2SO4.

174 g of K2SO4 yields 1520 g of FeSO4

8.7 g of K2SO4 will yield (1520/174)×8.7 = 76 g of FeSO4.

Hence, 76 g of Iron (II) sulphate is used .


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