NEON has 2 ISOTOPES of differing mass and in different proportions shown below. Which option demonstrates how we calculate its relative atomic mass (A_(r)) in g/mol {}^20Ne:91% {}^22Ne:9% ((91+9)times (20+22))/(100)=42 ((91times 9)+(20times 22))/(100)=12.6 ((91+20)times (9+22))/(100)=7.11 ((91times 20)+(9times 22))/(100)=20.2

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$NEON has 2 ISOTOPES of differing mass and in different proportions shown below. Which option demonstrates how we calculate its relative atomic mass (A_(r)) in g/mol {}^20Ne:91% {}^22Ne:9% ((91+9)times (20+22))/(100)=42 ((91times 9)+(20times 22))/(100)=12.6 ((91+20)times (9+22))/(100)=7.11 ((91times 20)+(9times 22))/(100)=20.2$

NEON has 2 ISOTOPES of differing mass and in different proportions shown below. Which option demonstrates how we calculate its relative atomic mass (A_(r)) in g/mol {}^20Ne:91% {}^22Ne:9% ((91+9)times (20+22))/(100)=42 ((91times 9)+(20times 22))/(100)=12.6 ((91+20)times (9+22))/(100)=7.11 ((91times 20)+(9times 22))/(100)=20.2

Answer 1

20.18 g/mol The relative atomic mass of Neon (Ne) is 20.18 g/mol. Among the options given, the correct calculation is $\frac{(91 \times 20)+(9 \times 22)}{100}=20.2$, which is approximately equal to our calculated value.

Answer 2

## Step 1: Understand the problem The problem is asking us to calculate the relative atomic mass of Neon (Ne). Neon has two isotopes, Ne-20 and Ne-22, with respective abundances of 91% and 9%. The relative atomic mass is calculated as the weighted average of the atomic masses of the isotopes, with the weights being the abundances of the isotopes. ## Step 2: Identify the correct formula The correct formula for calculating the relative atomic mass is to multiply the atomic mass of each isotope by its abundance (expressed as a decimal), and then sum up these results. ### Formula: $A = (m_1 \times p_1) + (m_2 \times p_2)$ where: - $m_1$ and $m_2$ are the atomic masses of the isotopes, - $p_1$ and $p_2$ are the abundances of the isotopes (expressed as decimals). ## Step 3: Substitute the values into the formula Substitute the atomic masses and abundances of the isotopes into the formula: ### $A = (20 \times 0.91) + (22 \times 0.09)$ ## Step 4: Perform the calculations Multiply the atomic masses by their respective abundances, and then add these results: ### $A = 18.2 + 1.98 = 20.18$

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