The three stable isotopes of neon: $^{20}_{10}\text{Ne}$, $^{21}_{10}\text{Ne}$ and $^{22}_{10}\text{Ne}$ have respective abundances of 90.51%, 0.27% and 9.22%. The atomic masses of the three isotopes are 19.99 u, 20.99 u and 21.99 u, respectively. Obtain the average atomic mass of neon.

The three stable isotopes of neon: ²⁰₁₀Ne, ²¹₁₀Ne and ²²₁₀Ne have respective abundances of 90.51%, 0.27% and 9

Exercises · 13.4

^{20}_{10}\text{Ne}

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✓ Free question

Multiply each isotopic mass by its fractional abundance and sum the three terms. The average atomic mass of neon comes out to $\approx 20.18\ \text{u}$ .

The weighted-average formula for three isotopes is:

M = f_1 m_1 + f_2 m_2 + f_3 m_3

Substituting the given abundances (as fractions) and masses:

M = (0.9051)(19.99\ \text{u}) + (0.0027)(20.99\ \text{u}) + (0.0922)(21.99\ \text{u})

Computing each term:

(0.9051)(19.99) = 18.0929\ \text{u}

(0.0027)(20.99) = 0.0567\ \text{u}

(0.0922)(21.99) = 2.0275\ \text{u}

Adding:

M = 18.0929 + 0.0567 + 2.0275 = 20.177\ \text{u} \approx 20.18\ \text{u}

This is very close to neon's accepted periodic-table atomic mass (20.18 u), confirming the calculation — the dominant isotope $^{20}\text{Ne}$ (90.5% abundant) pulls the average close to 20 u, with the heavier isotopes nudging it up slightly.

✓Final answer

M(\text{Ne}) \approx \boxed{20.18\ \text{u}}

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