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Atomic Weight Formula:
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Atomic weight (also called relative atomic mass) is the weighted average mass of an element's naturally occurring isotopes, taking into account their relative abundances. It's expressed in atomic mass units (amu).
The atomic weight is calculated using the formula:
Where:
Explanation: Multiply each isotope's mass by its natural abundance (as a decimal), then sum all these products to get the atomic weight.
Details: Atomic weight is fundamental in chemistry for stoichiometric calculations, determining molar masses, and understanding periodic trends. It's essential for laboratory work and industrial chemical processes.
Tips: Enter each isotope's mass in amu and its fractional abundance (e.g., 0.25 for 25%). You can add multiple isotopes. Ensure abundances sum to 1 (100%) for accurate results.
Q1: What's the difference between atomic weight and atomic mass?
A: Atomic mass refers to the mass of a single atom (usually of a specific isotope), while atomic weight is the weighted average of all naturally occurring isotopes.
Q2: Why do abundances need to be in decimal form?
A: The calculation requires fractional values (between 0 and 1) that represent proportions of the whole. For example, 25% abundance should be entered as 0.25.
Q3: What if my abundances don't sum to 1?
A: The calculator will still compute a result, but it won't accurately represent nature. For real elements, abundances should sum to 1 (100%).
Q4: How precise should the isotope masses be?
A: For most purposes, 4 decimal places are sufficient. High-precision work may require more decimal places.
Q5: Where can I find isotope mass and abundance data?
A: The IUPAC publishes authoritative data. Many chemistry textbooks and online databases also provide this information.