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Assoli18 [71]
3 years ago
10

How many grams of iron metal do you expect to be produced when 325 grams of an 87.5 percent by mass iron (II) nitrate solution r

eact with excess aluminum metal? Show all of the work needed to solve this problem.
2Al (s) + 3Fe(NO3)2 (aq) yields 3Fe (s) + 2Al(NO3)3 (aq
Chemistry
1 answer:
Basile [38]3 years ago
6 0
2Al (s) + 3Fe(NO3)2 (aq) → 3Fe (s) + 2Al(NO3)3 (aq)

Find the pure amount of iron nitrate solution:

325 g x 87.5% = 284.375 g

Convert 284.375g into mols by dividing by molar mass of the iron nitrate solution

Molar mass of iron nitrate =  55.85 g/mol + 2x14g/mol + 6x16g/mol = 179.85 g/mol

Moles of iron nitrate = 284.375 g / 179.85 g/mol = 1.58

Stoichiometry

3 mols of iron nitrate / 3 mols of iron = 1.58 mols of iron nitrate / x mols of iron

x = 1.58 mols of iron.

Convert that in grams by multiplying by the molar mass of iron

grams of iron = 1.58 mol x 55.85 g/mol = 88.24 g
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where,   e = 1.6 \times 10^{-19} C

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Hence,    U = \frac{1.423 \times 10^{-18} J}{1.6 \times 10^{-19} J}

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Therefore, U = 1.423 \times 10^{-18} J \times 1.67 \times 10^{-27} kJ/mol

                     = 2.37 \times 10^{-45} kJ/mol

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