An impure sample of zinc (zn) is treated with an excess of sulfuric acid (h 2 so 4) to form zinc sulfate (znso4) and molecular hydrogen (h 2). (a) write a balanced equation for the reaction. (b) if 0.0764 g of h 2 is obtained from 3.86 g of the sample, calculate the percent purity of the sample. (c) what assumptions must you make in (b)

a) The complete balanced chemical equation is:

Zn(s) + H2SO4(aq) > ZnSO4(aq) + H2 (g) 

b) First we find the amount zinc that has reacted based on the given H2 produced.

From stoichiometry 1 moles of Zn is needed for every 1 moles of H2 produced, therefore:

moles(Zn) = moles(H2) 

where moles is the ratio of mass and molar mass (MM)
mass(Zn) / MM(Zn) = mass(H2) / MM(H2) 
mass(Zn) = [mass(H2) / MM(H2)] * MM(Zn) 
mass(Zn) = [(0.0764 g)/(2 g/mol)] * 65.38 g/mol 
mass(Zn) = 2.49 g 

So we got 2.49 g pure zinc in the sample so the purity of zinc is therefore: 

purity = (2.49 / 3.86) * 100 % = 64.50 % 
 

c) In part (b) what we need to assume is that in the impurities in the sample do not react with sulfuric acid to produce hydrogen. So that the hydrogen is only from the reaction of Zn and sulphuric acid.

at the bottom

explanation:

m=9

step-by-step explanation: