A hypothetical element has two main isotopes with mass numbers of 56 and 59. If 61.00% of the isotopes have a mass number of 56 amu, what atomic weight should be listed on the periodic table for this element?
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To calculate the atomic weight of the element, you can use the average atomic mass formula, which takes into account the relative abundance of each isotope:
Atomic weight = (mass number₁ × abundance₁ + mass number₂ × abundance₂ + ...) / 100
Let's apply this formula to find the atomic weight of the element:
Let's assume that the isotope with a mass number of 56 amu has an abundance of 61.00%. The remaining abundance belongs to the isotope with a mass number of 59 amu.
Atomic weight = (56 × 61.00 + 59 × (100 - 61.00)) / 100
= (56 × 61.00 + 59 × 39.00) / 100
= (3416 + 2301) / 100
= 5717 / 100
= 57.17 amu
Therefore, the atomic weight that should be listed on the periodic table for this element is 57.17 amu.
To calculate the atomic weight for this element, we need to consider the mass numbers of the isotopes and the abundance of each isotope.
Given that 61.00% of the isotopes have a mass number of 56 amu, it means that the remaining 39.00% must have a mass number of 59 amu.
To calculate the atomic weight, we can use the following formula:
Atomic weight = (Abundance₁ * Mass₁) + (Abundance₂ * Mass₂)
Where:
- Abundance₁ and Abundance₂ are the percentages of abundances of the isotopes.
- Mass₁ and Mass₂ are the mass numbers of the isotopes.
Now, let's plug in the values:
Atomic weight = (0.6100 * 56 amu) + (0.3900 * 59 amu)
Calculating this equation will give you the atomic weight listed on the periodic table for this element.