Q: What is the percent composition by mass of oxygen in water ($\mathrm{H_2O}$)? (Atomic masses: H = 1.01, O = 16.00)

$\frac{16.00}{18.02} \times 100 = 88.8\%$

Q: Which of the following is the correct definition of a mole?

A mole is the amount of substance that contains $6.022 \times 10^{23}$ particles.

Q: What is the molar mass of ibuprofen ($\mathrm{C_{13}H_{18}O_2}$)? (Atomic masses: C = 12.01, H = 1.01, O = 16.00)

$13 \times 12.01 + 18 \times 1.01 + 2 \times 16.00 = 206.29$ g/mol

Q: If a reaction has a theoretical yield of 50.0 g and an actual yield of 45.8 g, what is the percent yield?

$\frac{45.8}{50.0} \times 100 = 91.6\%$

Q: How many hydrogen atoms are present in 10.0 g of water ($\mathrm{H_2O}$)? (Molar mass of water = 18.02 g/mol, Avogadro's number $N_A = 6.022 \times 10^{23}$)

$\frac{10.0}{18.02} \times 2 \times 6.022 \times 10^{23} = 6.69 \times 10^{23}$ atoms

Question 1

Which of the following statements best describes the relationship between atomic mass units (amu) and grams for a given element?

Accepted Answer

The mass of 1 atom in amu is numerically equal to the mass of 1 mole of that element in grams.

Question 2

What is the percent composition by mass of oxygen in water ($\mathrm{H_2O}$)? (Atomic masses: H = 1.01, O = 16.00)

Accepted Answer

$\frac{16.00}{18.02} 	imes 100 = 88.8\%$

Question 3

A sample contains 8.56 g of carbon and 1.44 g of hydrogen. What is the empirical formula of the compound? (Atomic masses: C = 12.01, H = 1.01)

Accepted Answer

$\mathrm{CH_4}$

Question 4

Which of the following is the correct definition of a mole?

Accepted Answer

A mole is the amount of substance that contains $6.022 	imes 10^{23}$ particles.

Question 5

How many moles are present in 0.568 g of iron (Fe)? (Atomic mass of Fe = 55.85 g/mol)

Accepted Answer

$\frac{0.568}{55.85} = 0.0102$ moles

Question 6

A chemist reacts 3.0 moles of nitrogen gas with 6.0 moles of hydrogen gas according to the equation $3\mathrm{H_2}(g) + \mathrm{N_2}(g) \rightarrow 2\mathrm{NH_3}(g)$. What is the limiting reactant?

Accepted Answer

Nitrogen gas ($\mathrm{N_2}$)

Question 7

If 5.22 g of NO reacts with 3.15 g of $\mathrm{O_2}$ according to $2\mathrm{NO}(g) + \mathrm{O_2}(g) \rightarrow 2\mathrm{NO_2}(g)$, what is the theoretical yield of $\mathrm{NO_2}$ in grams? (Molar masses: NO = 30.01 g/mol, $\mathrm{O_2}$ = 32.00 g/mol, $\mathrm{NO_2}$ = 46.01 g/mol)

Question 8

Which of the following is NOT a use of percent composition in real-world chemistry?

Accepted Answer

Calculating the number of moles in a sample

Question 9

What is the molar mass of ibuprofen ($\mathrm{C_{13}H_{18}O_2}$)? (Atomic masses: C = 12.01, H = 1.01, O = 16.00)

Accepted Answer

$13 	imes 12.01 + 18 	imes 1.01 + 2 	imes 16.00 = 206.29$ g/mol

Question 10

A compound has an empirical formula of $\mathrm{CF_2}$ and a molar mass of 200.0 g/mol. What is its molecular formula? (C = 12.01, F = 19.00)

Accepted Answer

$\mathrm{C_5F_{10}}$

Question 11

Which of the following best describes a limiting reactant in a chemical reaction?

Accepted Answer

The reactant that is completely consumed first, limiting the amount of product formed

Question 12

If a reaction has a theoretical yield of 50.0 g and an actual yield of 45.8 g, what is the percent yield?

Accepted Answer

$\frac{45.8}{50.0} 	imes 100 = 91.6\%$

Question 13

How many hydrogen atoms are present in 10.0 g of water ($\mathrm{H_2O}$)? (Molar mass of water = 18.02 g/mol, Avogadro's number $N_A = 6.022 	imes 10^{23}$)

Accepted Answer

$\frac{10.0}{18.02} 	imes 2 	imes 6.022 	imes 10^{23} = 6.69 	imes 10^{23}$ atoms

Question 14

Which conversion factor would you use to convert grams of a substance to moles?

Accepted Answer

Divide by molar mass

Question 15

A balanced chemical equation is required for stoichiometry because:

Accepted Answer

It provides the ratio of moles of reactants and products.

Question 16

If you have 10 slices of bread and 6 slices of cheese, and each sandwich requires 2 slices of bread and 1 slice of cheese, what is the maximum number of sandwiches you can make?

Accepted Answer

5

Question 17

Why is the atomic mass unit (amu) not practical for laboratory measurements?

Accepted Answer

Because laboratory balances measure mass in grams, not amu.

Question 18

Which of the following is the correct formula for calculating percent composition by mass of an element in a compound?

Accepted Answer

$\frac{	ext{mass of element in 1 mole}}{	ext{molar mass of compound}} 	imes 100$

The Mole Concept, Molar Mass, Percent Composition, and Stoichiometry

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