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Ch.2 - Atoms, Molecules, and Ions
Brown - Chemistry: The Central Science 15th Edition
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Tutti i libri di testoBrown 15th EditionCh.2 - Atoms, Molecules, and IonsProblema 35
Capitolo 2, Problema 35

Only two isotopes of copper occur naturally: 63Cu (atomic mass = 62.9296 amu; abundance 69.17%) 65Cu (atomic mass = 64.9278 amu; abundance 30.83%). Calculate the atomic weight (average atomic mass) of copper.

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1
Identify the atomic masses and their respective abundances for each isotope. For 63Cu, the atomic mass is 62.9296 amu and the abundance is 69.17%. For 65Cu, the atomic mass is 64.9278 amu and the abundance is 30.83%.
Convert the percentage abundances into decimal form by dividing each by 100. For 63Cu, 69.17% becomes 0.6917. For 65Cu, 30.83% becomes 0.3083.
Multiply the atomic mass of each isotope by its corresponding decimal abundance. For 63Cu, multiply 62.9296 amu by 0.6917. For 65Cu, multiply 64.9278 amu by 0.3083.
Add the products from the previous step to find the weighted average of the atomic masses. This sum will give you the average atomic mass of copper.
The result from the previous step is the atomic weight of copper, expressed in atomic mass units (amu).

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Isotopes

Isotopes are variants of a chemical element that have the same number of protons but different numbers of neutrons, resulting in different atomic masses. For example, copper has two stable isotopes, 63Cu and 65Cu, which differ in their neutron count and thus their atomic mass.
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Isotopes

Atomic Weight

Atomic weight, or average atomic mass, is the weighted average of the masses of an element's isotopes, taking into account their relative abundances. It is calculated by multiplying the mass of each isotope by its abundance (expressed as a fraction) and summing these values.
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Atom Structure

Weighted Average Calculation

To calculate a weighted average, each value is multiplied by its weight (or proportion), and the results are summed and divided by the total of the weights. In the case of copper, the atomic weight is calculated by summing the products of each isotope's mass and its abundance, providing a more accurate representation of the element's mass in nature.
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Average Rate Calculation Example
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Write the correct symbol, with both superscript and subscript, for each of the following. Use the list of elements in the front inside cover as needed: (d) the isotope of magnesium that has an equal number of protons and neutrons.

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(a) Thomson's cathode-ray tube (Figure 2.4) and the mass spectrometer (Figure 2.11) both involve the use of electric or magnetic fields to deflect charged particles. What are the charged particles involved in each of these experiments?

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Rubidium has two naturally occurring isotopes, rubidium-85 (atomic mass = 84.9118 amu; abundance = 72.15%) and rubidium-87 (atomic mass = 86.9092 amu; abundance = 27.85%). Calculate the atomic weight of rubidium

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Consider the mass spectrometer shown in Figure 2.11. Determine whether each of the following statements is true or false. If false, correct the statement to make it true: (a) The paths of neutral (uncharged) atoms are not affected by the magnet.

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The atomic weight of boron is reported as 10.81, yet no atom of boron has the mass of 10.81 amu. Which is the best explanation?

a. The measurement of atomic mass is only reliable to two significant figures.

b. The atomic weight is an average of many individual atoms.

c. The atomic weight is an average of many isotopes of the same nuclear composition.

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(b) Why is the atomic weight of carbon reported as 12.011 in the table of elements and the periodic table in the front inside cover of this text?

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