If a naturally occurring sample of an unidentified element is found to contain three isotopes (A, B, and C) and consists of 90.5% isotope A (mass number 20), 0.3% isotope B (mass number 21), and 9.2% isotope C (mass number 22), the atomic weight of the element measured from the sample will be:

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11ecba2d_124c_cb12_a3bf_3d763c51ba26_MD0008_00 Isotopes are atoms of the same element that have a different number of neutrons in the nucleus, and isotopes are identified by their mass number (the sum of the protons and neutrons). The masses of different isotopes are measured using the atomic mass unit (amu), which is defined as one-twelfth of the mass contained in a carbon-12 atom (the average mass of 6 protons and 6 neutrons). However, in a naturally occurring sample of an element, a mixture of isotopes is present in percentages resulting from their respective natural abundance. Because a naturally occurring sample of an element contains a mixture of isotopes, the atomic mass of a single isotope cannot be used in chemical calculations that involve bulk, macroscale samples. Instead, the atomic weight (more correctly called the relative atomic mass) is used, which accounts for the natural abundance of isotopes by taking a weighted contribution of the mass of each isotope found in the sample. It is these weighted averages that are listed as the atomic weights on the periodic table. For the unidentified element in this question, the overwhelming majority of the sample is isotope A, which has a mass number of 20. Therefore, the atomic weight (relative atomic mass) measured for the sample will be closest to 20 amu because isotope A is predominant within the sample. (Choices A and C) Isotopes B and C together account for only 9.5% of the sample. Therefore, the relative atomic mass cannot be greater than or equal to 21 amu. (Choice D) Although contributions from isotope B are small, they do have a slight impact on the value of the measured atomic weight. Educational objective: In a naturally occurring sample of an element, a mixture of isotopes is present in percentages resulting from their respective natural abundance. The atomic weights (relative atomic masses) listed on the periodic table account for natural isotope abundance by taking a weighted average of the mass of each isotope naturally found in a sample.

If a Naturally Occurring Sample of an Unidentified Element Is