An element has two naturally occurring isotopes. Isotope 1 has a mass of 120.9038 amu and a relative abundance of 57.4%, and isotope 2 has a mass of 122.9042 amu with a relative abundance of 42.6%. Calculate the atomic mass of this element and identify it.

• A. 121.00 amu, Antimony (Sb)
• B. 121.76 amu, Tin (Sn)
• C. 122.50 amu, Tellurium (Te)
• D. 120.50 amu, Indium (In)

Answer: (B)

Verified step by step guidance

1. Convert the relative abundances of the isotopes from percentages to decimal form by dividing each by 100. For Isotope 1, the relative abundance is 0.574, and for Isotope 2, it is 0.426.
2. Calculate the contribution of each isotope to the atomic mass by multiplying the mass of each isotope by its relative abundance. For Isotope 1, multiply 120.9038 amu by 0.574. For Isotope 2, multiply 122.9042 amu by 0.426.
3. Add the contributions from both isotopes to find the weighted average atomic mass of the element. This sum represents the atomic mass of the element.
4. Compare the calculated atomic mass to the given options to identify the element. The atomic mass should closely match one of the provided values.
5. Identify the element based on the calculated atomic mass and the options given. The element with an atomic mass closest to the calculated value is the correct identification.