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svlad2 [7]
3 years ago
11

Now we need to find the amount of NF3 that can be formed by the complete reactions of each of the reactants. If all of the N2 wa

s used up in the reaction, how many moles of NF3 would be produced
Chemistry
1 answer:
Alex777 [14]3 years ago
8 0

The question is incomplete, the complete question is:

Nitrogen and fluorine react to form nitrogen fluoride according to the chemical equation:

N_2(g)+3F_2(g)\rightarrow 2NF_3(g)

A sample contains 19.3 g of N_2 is reacted with 19.3 g of F_2. Now we need to find the amount of NF_3 that can be formed by the complete reactions of each of the reactants.

If all of the N_2 was used up in the reaction, how many moles of NF_3 would be produced?

<u>Answer:</u> 1.378 moles of NF_3 are produced in the reaction.

<u>Explanation:</u>

The number of moles is defined as the ratio of the mass of a substance to its molar mass.

\text{Number of moles}=\frac{\text{Given mass}}{\text{Molar mass}}       ......(1)

Limiting reagent is defined as the reagent which is completely consumed in the reaction and limits the formation of the product.

Excess reagent is defined as the reagent which is left behind after the completion of the reaction.

In the given chemical reaction, N_2 is considered as a limiting reagent because it limits the formation of the product and it was completely consumed in the reaction.

We are given:

Mass of N_2 = 19.3 g

Molar mass of N_2 = 28.02 g/mol

Putting values in equation 1:

\text{Moles of }N_2=\frac{19.3g}{28.02g/mol}=0.689mol

For the given chemical reaction:

N_2(g)+3F_2(g)\rightarrow 2NF_3(g)

By the stoichiometry of the reaction:

1 mole of N_2 produces 2 moles of NF_3

So, 0.689 moles of N_2 will produce = \frac{2}{1}\times 0.689=1.378mol of NF_3

Hence, 1.378 moles of NF_3 are produced in the reaction.

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