Q: Which equation correctly calculates the mass percent of chlorine in $$\mathrm{CCl}_2\mathrm{F}_2$$?

$$\dfrac{2 \times \text{molar mass of Cl}}{\text{molar mass of } \mathrm{CCl}_2\mathrm{F}_2} \times 100\%$$

Q: Which of the following is the correct formula to determine the value of $$n$$ for a molecular formula?

$$n = \dfrac{\text{molar mass}}{\text{empirical formula molar mass}}$$

Question 1

Which of the following best describes the empirical formula of a compound?

Accepted Answer

It is the simplest whole-number ratio of atoms of each element in a compound.

Question 2

A compound is found to contain 24.5 g nitrogen and 70.0 g oxygen. What is the empirical formula of the compound? (Atomic masses: N = 14.01 g/mol, O = 16.00 g/mol)

Accepted Answer

$$\mathrm{NO}_2$$

Question 3

Which equation correctly calculates the mass percent of chlorine in $$\mathrm{CCl}_2\mathrm{F}_2$$?

Accepted Answer

$$\dfrac{2 	imes 	ext{molar mass of Cl}}{	ext{molar mass of } \mathrm{CCl}_2\mathrm{F}_2} 	imes 100\%$$

Question 4

A sample of water is decomposed in the laboratory to produce 2.02 g of hydrogen and 16.00 g of oxygen. What is the empirical formula of water?

Accepted Answer

$$\mathrm{H}_2\mathrm{O}$$

Question 5

A compound used to purify drinking water contains 15.8% sodium, 28.1% sulfur, and 56.1% oxygen by mass. What is its empirical formula?

Accepted Answer

$$\mathrm{Na}_2\mathrm{SO}_4$$

Question 6

The empirical formula of fructose is $$\mathrm{CH}_2\mathrm{O}$$ and its molar mass is 180.2 g/mol. What is the molecular formula of fructose?

Accepted Answer

$$\mathrm{C}_6\mathrm{H}_{12}\mathrm{O}_6$$

Question 7

Which of the following is the correct formula to determine the value of $$n$$ for a molecular formula?

Accepted Answer

$$n = \dfrac{	ext{molar mass}}{	ext{empirical formula molar mass}}$$

Question 8

During combustion analysis, a compound containing only carbon and hydrogen produces 1.83 g $$\mathrm{CO}_2$$ and 0.901 g $$\mathrm{H}_2\mathrm{O}. $$What is the empirical formula of the compound? (Molar masses: $$\mathrm{CO}_2 = 44.01 g/$$mol, $$\mathrm{H}_2\mathrm{O} = 18.02 g/$$mol)

Accepted Answer

$$\mathrm{CH}_2$$

Question 9

Which statement about balancing chemical equations is correct?

Accepted Answer

Coefficients are changed to balance the equation, but subscripts should never be changed.

Question 10

What is the balanced chemical equation for the combustion of butane ($$\mathrm{C}_4\mathrm{H}_{10}$$)?

Accepted Answer

$$2\mathrm{C}_4\mathrm{H}_{10}(g) + 13\mathrm{O}_2(g) \rightarrow 8\mathrm{CO}_2(g) + 10\mathrm{H}_2\mathrm{O}(g)$$

Question 11

In a combustion analysis, how is the amount of oxygen in a compound determined?

Accepted Answer

By subtracting the masses of carbon and hydrogen from the original mass of the sample.

Question 12

Which of the following is a correct symbolic representation of a balanced chemical reaction for the combustion of methane ($$\mathrm{CH}_4$$)?

Accepted Answer

$$\mathrm{CH}_4 + 2\mathrm{O}_2 \rightarrow \mathrm{CO}_2 + 2\mathrm{H}_2\mathrm{O}$$

Question 13

If a chemical equation contains a fractional coefficient, what must be done to obtain the final balanced equation?

Accepted Answer

Multiply all coefficients by the denominator of the fraction to obtain whole numbers.

Question 14

Which of the following best describes the difference between subscripts and coefficients in a chemical equation?

Accepted Answer

Subscripts indicate the number of atoms in a molecule; coefficients indicate the number of molecules.

Question 15

A compound contains 40.0% carbon, 6.7% hydrogen, and 53.3% oxygen by mass. What is its empirical formula? (Atomic masses: C = 12.01 g/mol, H = 1.01 g/mol, O = 16.00 g/mol)

Accepted Answer

$$\mathrm{CH}_2\mathrm{O}$$

Question 16

Which of the following is NOT a correct step in determining an empirical formula from mass percent data?

Accepted Answer

Change the subscripts in the chemical formula to balance the equation.

Question 17

Which statement about chemical reactions is correct?

Accepted Answer

Atoms are rearranged during a chemical reaction, but are not created or destroyed.

General Chemistry I: Composition of Compounds and Chemical Reactions

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