I see that you've uploaded a page with several problems related to calculating the average atomic mass of elements such as Argon (Ar), Potassium (K), and Calcium (Ca). The average atomic mass is calculated using the formula:

$\text{Average Atomic Mass} = \sum (\text{Mass of Isotope} \times \text{Abundance Fraction})$

Let's break down one of these calculations. I will start with Argon (Ar):

Problem 18: Calculate the Average Atomic Mass of Argon (Ar)

The isotopes and their respective data are:

• $^{36}\text{Ar}$: Mass = 35.967546 amu, Abundance = 0.3365%
• $^{38}\text{Ar}$: Mass = 37.962732 amu, Abundance = 0.0632%
• $^{40}\text{Ar}$: Mass = 39.962383 amu, Abundance = 99.6003%

Convert the percentage abundances to fractions by dividing by 100:

• $0.3365\% = 0.003365$
• $0.0632\% = 0.000632$
• $99.6003\% = 0.996003$

Now, apply the formula: $\text{Average Atomic Mass} = (35.967546 \times 0.003365) + (37.962732 \times 0.000632) + (39.962383 \times 0.996003)$

I will calculate this now.The calculated average atomic mass of Argon (Ar) is approximately 39.9477 amu.

Would you like me to proceed with the calculations for Potassium (K) or Calcium (Ca), or do you have any specific questions?

Here are five related questions for further exploration:

1. How do isotopic abundances influence the periodic table's listed atomic mass?
2. What are the most common isotopes for naturally occurring elements?
3. How do we experimentally determine isotopic abundances?
4. Why is the average atomic mass not always a whole number?
5. How does the average atomic mass vary across different samples of the same element?

Tip: Always convert percentages into fractions (divide by 100) when using them in weighted average calculations.