02 The Mole and Chemical Composition Online Quiz Questions

Using $12.01\ \mathrm{g/mol}$ for carbon and $1.008\ \mathrm{g/mol}$ for hydrogen, calculate the molar mass of methane, $\mathrm{CH_4}$, and the mass percentage of each element. Show the mass contribution used for each percentage and check that the percentages total approximately $100\%$.

A $29.22\ \mathrm{g}$ sample of NaCl has a molar mass of $58.44\ \mathrm{g/mol}$. Determine the amount of NaCl in moles and the number of NaCl formula units in the sample. Use $6.02214076\times10^{23}$ entities per mole, and report the entity type explicitly.

A sample contains $0.125\ \mathrm{mol}$ of $\mathrm{H_2O}$. Use $6.02214076\times10^{23}$ entities per mole to determine the number of water molecules, hydrogen atoms, and oxygen atoms. Report each count to three significant figures and label what is being counted.

A compound contains $24.02\ \mathrm{g}$ C, $3.024\ \mathrm{g}$ H, and $16.00\ \mathrm{g}$ O. Using molar masses of $12.01$, $1.008$, and $16.00\ \mathrm{g/mol}$, respectively, determine the empirical formula. Show the mole ratio and how you reduce it.

A compound is $52.2\%$ C, $13.0\%$ H, and $34.8\%$ O by mass. Using molar masses of $12.01$, $1.008$, and $16.00\ \mathrm{g/mol}$, respectively, determine its empirical formula. Show how you convert the mass percentages to a simplest whole-number atom ratio.

After converting a compound’s elemental masses to moles and dividing by the smallest amount, you obtain $\mathrm{C:H:O}=1.00:1.48:1.00$. Determine the most reasonable empirical formula and explain why simply rounding $1.48$ to $1$ would be a poor choice.

An oxide is $27.3\%$ C and $72.7\%$ O by mass. Using molar masses of $12.01\ \mathrm{g/mol}$ for C and $16.00\ \mathrm{g/mol}$ for O, determine its empirical formula. Explain why the two percentages themselves cannot be used directly as the atom ratio.

A compound has empirical formula $\mathrm{C_2H_3O}$ and molecular molar mass $129.13\ \mathrm{g/mol}$. Using $12.01$, $1.008$, and $16.00\ \mathrm{g/mol}$ for C, H, and O, respectively, determine its molecular formula. Show the empirical-formula mass and multiplier.

Two compounds share the empirical formula $\mathrm{C_2H_3O}$, but their molecular molar masses are $86.09\ \mathrm{g/mol}$ and $129.13\ \mathrm{g/mol}$. Using atomic molar masses C $=12.01$, H $=1.008$, and O $=16.00\ \mathrm{g/mol}$, determine the molecular formula for each compound and explain what this comparison shows about empirical formulas.

A sample is reported to have empirical formula $\mathrm{C_2H_3O}$ and molecular molar mass $77.0\ \mathrm{g/mol}$. Using C $=12.01$, H $=1.008$, and O $=16.00\ \mathrm{g/mol}$, assess whether these data can yield a valid molecular formula. Show the calculation and state the appropriate conclusion.

An unknown compound has empirical formula $\mathrm{C_3H_4O}$ and molecular molar mass $168.19\ \mathrm{g/mol}$. Use atomic molar masses C $=12.01$, H $=1.008$, and O $=16.00\ \mathrm{g/mol}$ to determine the molecular formula. Show the empirical-formula mass and demonstrate that the result is a whole-number multiple of the empirical formula.

Calcium hydroxide has formula $\mathrm{Ca(OH)_2}$. Using atomic molar masses of $40.08\ \mathrm{g/mol}$ for Ca, $16.00\ \mathrm{g/mol}$ for O, and $1.008\ \mathrm{g/mol}$ for H, calculate the mass of a $3.20$-mol sample. Report your result to three significant figures.

A KBr sample has a mass of $11.90\ \mathrm{g}$, and its molar mass is $119.00\ \mathrm{g/mol}$. How many KBr formula units does the sample contain? Use $N_A = 6.02214076 \times 10^{23}\ \mathrm{entities/mol}$, and report your result to four significant figures.

For $0.250\ \mathrm{mol}$ of $\mathrm{Al_2(SO_4)_3}$ formula units, determine the amount in moles of Al atoms, S atoms, and O atoms separately, then find the total amount of atoms.

Calculate the mass percent of each element in $\mathrm{Na_2SO_4}$. Use atomic molar masses of $22.99\ \mathrm{g/mol}$ for Na, $32.06\ \mathrm{g/mol}$ for S, and $16.00\ \mathrm{g/mol}$ for O. Show the molar mass and the three percentages.

A compound is $27.3\%$ carbon and $72.7\%$ oxygen by mass. Using atomic molar masses of $12.01\ \mathrm{g/mol}$ for C and $16.00\ \mathrm{g/mol}$ for O, determine its empirical formula. Show how you obtain the simplest whole-number ratio.

A sample contains $6.01\ \mathrm{g}$ C, $0.756\ \mathrm{g}$ H, and $8.00\ \mathrm{g}$ O. Using atomic molar masses of $12.01\ \mathrm{g/mol}$ for C, $1.008\ \mathrm{g/mol}$ for H, and $16.00\ \mathrm{g/mol}$ for O, determine the empirical formula. Show how you handle any non-whole-number ratio.

A compound has empirical formula $\mathrm{CH_2}$ and molecular molar mass $84.16\ \mathrm{g/mol}$. Using atomic molar masses of $12.01\ \mathrm{g/mol}$ for C and $1.008\ \mathrm{g/mol}$ for H, determine its molecular formula. Show the multiplier.

A compound has empirical formula $\mathrm{C_2H_6O}$ and molecular molar mass $92.14\ \mathrm{g/mol}$. Using atomic molar masses of $12.01\ \mathrm{g/mol}$ for C, $1.008\ \mathrm{g/mol}$ for H, and $16.00\ \mathrm{g/mol}$ for O, determine its molecular formula and explain how each subscript changes.

A compound is $30.45\%$ N and $69.55\%$ O by mass, and its molecular molar mass is $92.02\ \mathrm{g/mol}$. Using atomic molar masses of $14.01\ \mathrm{g/mol}$ for N and $16.00\ \mathrm{g/mol}$ for O, determine both its empirical formula and its molecular formula. Show the calculation that links the two.

Explain why an empirical formula alone cannot determine a compound's molecular formula. Give two different molecular formulas with the same empirical formula, and identify what additional information can determine the multiplier.

A compound is reported to have an empirical-formula mass of $30.03\ \mathrm{g/mol}$ and a molecular molar mass of $75.0\ \mathrm{g/mol}$. Calculate the implied multiplier and assess whether these values support a molecular formula under the whole-number-multiple rule. Explain what should be done rather than rounding the multiplier without justification.

An analysis reports mass percentages of $24.7\%$, $18.6\%$, and $56.5\%$ for the elements in a compound. Find their total and explain whether the result, by itself, means the composition is invalid.

A sample contains $0.250\ \mathrm{mol}$ of $\mathrm{H_2O}$. Calculate the number of water molecules, hydrogen atoms, and oxygen atoms in the sample. Report entity counts to three significant figures, and identify what each count refers to.

A laboratory compares $1.25\ \mathrm{mol}$ samples of NaCl and $\mathrm{H_2O}$. Using molar masses of $58.44\ \mathrm{g/mol}$ for NaCl and $18.02\ \mathrm{g/mol}$ for $\mathrm{H_2O}$, calculate each sample's mass to three significant figures. Explain why equal amounts in moles do not have equal masses.

Using atomic molar masses $\mathrm{C}=12.01$, $\mathrm{H}=1.008$, and $\mathrm{O}=16.00\ \mathrm{g/mol}$, calculate the mass percent of each element in $\mathrm{C_2H_6O}$. Show how you obtain the molar mass and each element's contribution, and check the total percentage.

A sample contains $24.02\ \mathrm{g}$ of carbon, $3.024\ \mathrm{g}$ of hydrogen, and $32.00\ \mathrm{g}$ of oxygen. Using atomic molar masses $\mathrm{C}=12.01$, $\mathrm{H}=1.008$, and $\mathrm{O}=16.00\ \mathrm{g/mol}$, determine the empirical formula. Show the mole ratio and how you reduce it.

A compound has empirical formula $\mathrm{CH_2O}$ and molecular molar mass $180.16\ \mathrm{g/mol}$. Using atomic molar masses $\mathrm{C}=12.01$, $\mathrm{H}=1.008$, and $\mathrm{O}=16.00\ \mathrm{g/mol}$, determine its molecular formula. Show the empirical-formula mass, the multiplier, and how the subscripts change.

Experimental mole amounts for carbon, hydrogen, and oxygen produce the ratio $1.00:1.50:2.00$. Determine the simplest whole-number ratio and corresponding empirical formula. Explain why rounding $1.50$ to the nearest whole number before reducing the ratio would be a poor choice.

Ethanol is written as $\mathrm{C_2H_5OH}$. Using $12.01\ \mathrm{g/mol}$ for carbon, $1.008\ \mathrm{g/mol}$ for hydrogen, and $16.00\ \mathrm{g/mol}$ for oxygen, calculate its molar mass. Show the sum of the atomic contributions and explain how many hydrogen atoms are included in that sum.