Atomic Mass Calculator

How to Calculate Atomic Mass from Isotopes

Atomic Mass Formula:

\[ \text{Atomic Mass} = \sum (\text{Isotope Mass} \times \text{Abundance}) \]

Isotope 1

Isotope Mass:

Natural Abundance:

Calculated Atomic Mass:

Unit Converter ▲

Unit Converter ▼

From:	To:

1. What is Atomic Mass?

Atomic mass (also called relative atomic mass) is the weighted average mass of all naturally occurring isotopes of an element, taking into account their natural abundances. It's expressed in atomic mass units (u), where 1 u is 1/12 the mass of a carbon-12 atom.

2. How to Calculate Atomic Mass

The atomic mass is calculated using the formula:

\[ \text{Atomic Mass} = \sum (\text{Isotope Mass} \times \text{Abundance}) \]

Where:

Isotope Mass — Mass of each isotope in atomic mass units (u)
Abundance — Natural abundance of each isotope (as a decimal)

Explanation: Multiply each isotope's mass by its natural abundance (as a decimal), then sum all these values to get the weighted average atomic mass.

3. Importance of Atomic Mass

Details: Atomic mass is fundamental in chemistry for stoichiometric calculations, determining molar masses, and understanding chemical reactions. It appears on the periodic table for each element.

4. Using the Calculator

Tips:

Enter the mass of each isotope in atomic mass units (u)
Enter the natural abundance as a percentage (e.g., 75.77 for 75.77%)
Click "Add Another Isotope" for elements with multiple isotopes
Click "Calculate" to compute the weighted average atomic mass

5. Frequently Asked Questions (FAQ)

Q1: Why is atomic mass not a whole number?
A: Atomic mass is a weighted average of all naturally occurring isotopes, which rarely results in a whole number.

Q2: What's the difference between mass number and atomic mass?
A: Mass number is the sum of protons and neutrons in a specific isotope (always a whole number), while atomic mass is the weighted average of all isotopes.

Q3: How many isotopes should I include?
A: Include all naturally occurring isotopes with significant abundance (>0.1%). For most elements, 1-5 isotopes are sufficient.

Q4: Why does my calculation differ slightly from the periodic table?
A: The periodic table uses more precise measurements and may include very rare isotopes not accounted for in your calculation.

Q5: Can I use this for radioactive elements?
A: For radioactive elements with very long half-lives (like uranium), yes. For short-lived radioisotopes, the concept of "natural abundance" doesn't apply.