Question

A sample of salt has the following percentage composition:        Fe= 36.76 %       S= 21.11%   O= 42.14 %     Calculate the empirical formula of the compound. What happens when  the salt is heated?  Write a balanced chemical equation. (atomic mass of Fe=56 u , S=32 u, O=16 u)​

Answer 1

Answer:

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Answer 2

Given:

Percentage composition: Fe= 36.76 %  , S= 21.11%   ,  O= 42.14 %

To find:

Empirical formula of the compound

Explanation :

• Assuming that 100 g of the compound is present and further changing the percent to grams:

Fe →  36.76 g

S   →   21.11 g

O  →   42.14 g

• Masses are converted to moles

Fe →    36.76/55.84 = 0.65

S   →    21.11/32 = 0.65

O  →    42.14/16 = 2.63

• Dividing by the lowest integer, converting into the smallest whole-number ratio

Fe  →  0.65/0.65 = 1

S    →  0.65/0.65= 1

O   →   2.63/0.65 = 4

• Therefore the empirical formula of the compound is $FeSO_4$ .

• Action of heat on the salt , $FeSO_4$

$2FeSO_4 \rightarrow Fe_2O_3 + SO_2 + SO_3$

• When heated, iron(II) sulphate loses all the water of crystallization and green color are converted into a dirty-yellow anhydrous solid. With continued heating , a reddish brown iron oxide is formed along with the release of $SO_2$ and white fumes of $SO_3$.