Question

A certain element consist of two stable isotopes . The first has a mass of 14.0031 amu and a percent natural abundance of 99.63% . The second has mass of 15.001 amu and a percent natural abundance of 0.37% . What is the atomic weight of the element

Answer 1

Answer:

13.98886 atomic mass units (amu).

Explanation:

To find the atomic weight of the element, we can use the formula:

Atomic weight = (mass1 * abundance1 + mass2 * abundance2) / 100

Where:

mass1 and mass2 are the masses of the isotopes

abundance1 and abundance2 are the percent natural abundances of the isotopes

Given:

mass1 = 14.0031 amu

abundance1 = 99.63%

mass2 = 15.001 amu

abundance2 = 0.37%

Let's plug these values into the formula:

Atomic weight = (14.0031 * 99.63 + 15.001 * 0.37) / 100

≈ (1393.3333 + 5.55337) / 100

≈ 1398.8863 / 100

≈ 13.98886 amu

So, the atomic weight of the element is approximately 13.98886 atomic mass units (amu).