Q: Given two isotopes of copper: Cu-63 (mass = 62.9296 amu, abundance = 69.17%) and Cu-65 (mass = 64.9278 amu, abundance = 30.83%), what is the atomic mass of copper?

$$62.9296 \times 0.6917 + 64.9278 \times 0.3083 = 63.55$$ amu

Q: Which of the following best describes Avogadro’s number?

The number of particles in 1 mole, $$6.022 \times 10$$^{23}$$.

Q: A silver sphere contains $$3.76 \times 10$$^{22}$$ atoms. If the molar mass of silver is 107.87$ g/mol and its density is 10.49$ g/cm^3$, what is the radius of the sphere in cm?

First, calculate the mass: $$\frac{3.76 \times 10$$^{22}$$}{6.022 \times 10$$^{23}$$} \times 107.87 g. $$Then, volume: $$\frac{\text{mass}}{10.49} cm$^3. $$Finally, $$r = \left(\frac{3V}{4\pi}\$$right)^{1/3}$$. The radius is approximately 0.62$ cm.

Q: Which formula correctly calculates the average atomic mass of an element with two isotopes?

$$\text{Atomic mass} = (\text{abundance}_1 \times \text{mass}_1) + (\text{abundance}_2 \times \text{mass}_2)$$

Q: If you have 12.0 g of carbon (atomic mass = 12.01 amu), how many atoms of carbon do you have?

$$\frac{12.0}{12.01} \times 6.022 \times 10$$^{23}$$ = 6.02 \times 10$$^{23}$$ atoms

Q: A sample contains $$2.0 cm $$radius sphere of indium (density $$7.31 g/cm$^3). If $$natural indium is $$4.3\% In-113 $$and $$95.7\% In-115$$, how many neutrons are present in the sphere?

Calculate volume: $$V = \frac{4}{3}\pi r^3$. Mass: m = V$$ \times 7.31$$. Moles: $$\frac{m}{114.82}$$ g/mol. Atoms: $$\text{moles} \times 6.022 \times $$10^{23}. $$Neutrons: $$0.043 \times \text{atoms} \times 66 + 0.957 \times \text{atoms} \times 68.$$

Q: If a sample contains $$1.50$$ moles of aluminum atoms, what is its mass in grams? (Molar mass of Al = $$26.98 g/$$mol)

$$1.50 \times 26.98 = 40.47 g$$

Question 1

Which statement best explains why the atomic mass of chlorine is not a whole number?

Accepted Answer

Because the atomic mass is an average of the masses of all naturally occurring isotopes, weighted by their natural abundance.

Question 2

Given two isotopes of copper: Cu-63 (mass = 62.9296 amu, abundance = 69.17%) and Cu-65 (mass = 64.9278 amu, abundance = 30.83%), what is the atomic mass of copper?

Accepted Answer

$$62.9296 	imes 0.6917 + 64.9278 	imes 0.3083 = 63.55$$ amu

Question 3

Which of the following best describes Avogadro’s number?

Accepted Answer

The number of particles in 1 mole, $$6.022 	imes 10$$^{23}$$.

Question 4

A silver sphere contains $$3.76 	imes 10$$^{22}$$ atoms. If the molar mass of silver is 107.87$ g/mol and its density is 10.49$ g/cm^3$, what is the radius of the sphere in cm?

Accepted Answer

First, calculate the mass: $$\frac{3.76 	imes 10$$^{22}$$}{6.022 	imes 10$$^{23}$$} 	imes 107.87 g. $$Then, volume: $$\frac{	ext{mass}}{10.49} cm$^3. $$Finally, $$r = \left(\frac{3V}{4\pi}\$$right)^{1/3}$$. The radius is approximately 0.62$ cm.

Question 5

Which formula correctly calculates the average atomic mass of an element with two isotopes?

Accepted Answer

$$	ext{Atomic mass} = (	ext{abundance}_1 	imes 	ext{mass}_1) + (	ext{abundance}_2 	imes 	ext{mass}_2)$$

Question 6

If you have 12.0 g of carbon (atomic mass = 12.01 amu), how many atoms of carbon do you have?

Accepted Answer

$$\frac{12.0}{12.01} 	imes 6.022 	imes 10$$^{23}$$ = 6.02 	imes 10$$^{23}$$ atoms

Question 7

Which statement about isotopes is correct?

Accepted Answer

Isotopes of an element have the same number of protons but different numbers of neutrons.

Question 8

A sample contains $$2.0 cm $$radius sphere of indium (density $$7.31 g/cm$^3). If $$natural indium is $$4.3\% In-113 $$and $$95.7\% In-115$$, how many neutrons are present in the sphere?

Accepted Answer

Calculate volume: $$V = \frac{4}{3}\pi r^3$. Mass: m = V$$ 	imes 7.31$$. Moles: $$\frac{m}{114.82}$$ g/mol. Atoms: $$	ext{moles} 	imes 6.022 	imes $$10^{23}. $$Neutrons: $$0.043 	imes 	ext{atoms} 	imes 66 + 0.957 	imes 	ext{atoms} 	imes 68.$$

Question 9

Which of the following is the correct definition of molar mass?

Accepted Answer

The mass of one mole of atoms of an element, numerically equal to the atomic mass in amu but expressed in grams.

Question 10

Which experimental technique is most commonly used to determine the natural abundance of isotopes?

Accepted Answer

Mass spectrometry

Question 11

If a sample contains $$1.50$$ moles of aluminum atoms, what is its mass in grams? (Molar mass of Al = $$26.98 g/$$mol)

Accepted Answer

$$1.50 	imes 26.98 = 40.47 g$$

Question 12

Which of the following statements about the mole is correct?

Accepted Answer

A mole is $$6.022 	imes 10$$^{23}$$ particles of any substance.

Question 13

Which calculation would you use to find the number of atoms in 5.00 g of helium (molar mass = 4.00 g/mol)?

Accepted Answer

$$\frac{5.00}{4.00} 	imes 6.022 	imes 10$$^{23}$$ atoms

Question 14

Which of the following best describes the relationship between atomic mass (amu) and molar mass (g/mol)?

Accepted Answer

Atomic mass in amu is numerically equal to the molar mass in g/mol for an element.

Question 15

A chemist needs to prepare $$1.00 	imes 10$$^{24}$$ atoms of carbon. What mass of carbon should be weighed out? (Atomic mass = 12.01 amu)

Accepted Answer

$$\frac{1.00 	imes 10$$^{24}$$}{6.022 	imes 10$$^{23}$$} 	imes 12.01 = 19.96 g$$

Question 16

Which of the following is NOT a correct step in calculating the percent abundance of an isotope from mass spectrometry data?

Accepted Answer

Add the atomic numbers of all isotopes together.

Atoms and Elements: Atomic Mass, Isotopes, and Counting Atoms

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