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# Mass-Mass Stoichiometry

## Calculations for converting between masses of different reactants and products

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Practice Mass-Mass Stoichiometry
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Mass-Mass Stoichiometry

Credit: Pauline Mak
Source: http://www.flickr.com/photos/__my__photos/5591677002/

#### How many walnuts are needed to equal 250 grams?

I want to send 250 grams of shelled walnuts to a friend (don’t ask why – just go with the question). How many walnuts in shells do I need to buy? To figure this out, I need to know how much the shell of a walnut weighs (about 40% of the total weight of the unshelled walnut). I can then calculate the mass of walnuts that will give me 250 grams of shelled walnuts and then determine how many walnuts I need to buy.

### Mass to Mass Problems

Mass-mass calculations are the most practical of all mass-based stoichiometry problems. Moles cannot be measured directly, while the mass of any substance can generally be easily measured in the lab. This type of problem is three steps and is a combination of the two previous types.

mass of givenmoles of givenmoles of unknownmass of unknown\begin{align*}\text{mass of} \ given \rightarrow \text{moles of} \ given \rightarrow \text{moles of} \ unknown \rightarrow \text{mass of} \ unknown\end{align*}

The mass of the given substance is converted into moles by use of the molar mass of that substance from the periodic table. Then, the moles of the given substance are converted into moles of the unknown by using the mole ratio from the balanced chemical equation. Finally, the moles of the unknown are converted to mass by use of its molar mass.

#### Sample Problem: Mass-Mass Stoichiometry

Ammonium nitrate decomposes to dinitrogen monoxide and water according to the following equation.

NH4NO3(s)N2O(g)+2H2O(l)\begin{align*}\text{NH}_4\text{NO}_3 (s) \rightarrow \text{N}_2 \text{O}(g)+2\text{H}_2\text{O}(l)\end{align*}

In a certain experiment, 45.7 g of ammonium nitrate is decomposed. Find the mass of each of the products formed.

Step 1: List the known quantities and plan the problem.

Known

• given: 45.7 g NH4NO3
• 1 mol NH4NO3 = 1 mol N2O = 2 mol H2O (mole ratios)
• molar mass of NH4NO3 = 80.06 g/mol
• molar mass of N2O = 44.02 g/mol
• molar mass of H2O = 18.02 g/mol

Unknown

• mass N2O = ? g
• mass H2O = ? g

Perform two separate three-step mass-mass calculations as shown below.

g NH4NO3mol NH4NO3mol N2Og N2Og NH4NO3mol NH4NO3mol H2Og H2O\begin{align*}& \text{g NH}_4\text{NO}_3 \rightarrow \text{mol NH}_4\text{NO}_3 \rightarrow \text{mol N}_2\text{O} \rightarrow \text{g N}_2\text{O}\\ & \text{g NH}_4\text{NO}_3 \rightarrow \text{mol NH}_4\text{NO}_3 \rightarrow \text{mol H}_2\text{O} \rightarrow \text{g H}_2\text{O}\end{align*}

Step 2: Solve.

45.7 g NH4NO3×1 mol NH4NO380.06 g NH4NO3×1 mol N2O1 mol NH4NO3×44.02 g N2O1 mol N2O=25.1 g N2O45.7 g NH4NO3×1 mol NH4NO380.06 g NH4 NO3×2 mol H2O1 mol NH4NO3×18.02 g H2O1 mol H2O=20.6 g N2O\begin{align*}& 45.7 \text{ g NH}_4\text{NO}_3 \times \frac{1 \text{ mol NH}_4\text{NO}_3}{80.06 \text{ g NH}_4\text{NO}_3} \times \frac{1 \text{ mol N}_2 \text{O}}{1 \text{ mol NH}_4\text{NO}_3} \times \frac{44.02 \text{ g N}_2\text{O}}{1 \text{ mol N}_2\text{O}}=25.1 \text{ g N}_2\text{O}\\ & 45.7 \text{ g NH}_4\text{NO}_3 \times \frac{1 \text{ mol NH}_4\text{NO}_3}{80.06 \text{ g NH}_4 \ \text{NO}_3} \times \frac{2 \text{ mol H}_2\text{O}}{1 \text{ mol NH}_4\text{NO}_3} \times \frac{18.02 \text{ g H}_2\text{O}}{1 \text{ mol H}_2\text{O}}=20.6 \text{ g N}_2\text{O}\end{align*}

The total mass of the two products is equal to the mass of ammonium nitrate which decomposed, demonstrating the law of conservation of mass. Each answer has three significant figures.

### Summary

• Mass-mass calculations involve converting the mass of a reactant to moles of reactant, then using mole ratios to determine moles of product which can then be converted to mass of product.

### Review

1. If matter is neither created nor destroyed, why can’t we just go directly from grams of reactant to grams of product?
2. Why is it important to get the subscripts correct in the formulas?
3. Why do the coefficients need to be correct?

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