Answer: Atomic masses are not whole numbers because they take into account the presence of isotopes, which have different masses, and the atomic mass listed on the periodic table is a weighted average of these isotopic masses.
The atomic masses of elements are not whole numbers because they consider the existence of isotopes within a given element. Isotopes are atoms of the same element that have the same number of protons but different numbers of neutrons, resulting in different atomic masses.
Here’s a detailed explanation:
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Isotopes and Atomic Mass:
- Isotopes of an element have the same atomic number (number of protons) but different mass numbers (sum of protons and neutrons).
- The atomic mass of an element is the weighted average of the masses of its isotopes, taking into account the abundance of each isotope in nature.
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Weighted Average Calculation:
- The atomic mass listed on the periodic table is not a simple average but a weighted average.
- Each isotope’s contribution to the atomic mass is weighted by its abundance, which is the percentage of that isotope found in a natural sample of the element.
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Fractional Abundance:
- Isotopes are typically found in nature in varying proportions, and these proportions are expressed as fractional abundances.
- When calculating the atomic mass, these fractional abundances result in a weighted sum that may not yield a whole number.
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Example with Carbon:
- Carbon has two stable isotopes: carbon-12 and carbon-13.
- Carbon-12 is more abundant in nature, and carbon-13 is less abundant.
- The atomic mass of carbon is a weighted average, considering the abundance of each isotope, and is approximately 12.01 atomic mass units (u).
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Variation Across the Periodic Table:
- Different elements may have a greater or lesser variation in their isotopic masses and abundances.
- Elements with only one stable isotope or isotopes with similar masses and abundances may have atomic masses that are closer to whole numbers.
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Technological Advances:
- Advances in mass spectrometry and other analytical techniques have allowed for more accurate measurements of isotopic abundances and, consequently, more precise determination of atomic masses.
In summary, atomic masses are not whole numbers because they reflect the weighted average of the masses of an element’s isotopes, considering their fractional abundances in nature. This approach provides a more accurate representation of the actual atomic mass of an element as observed in natural samples.