Question
How do I calculate atomic mass?
Answers
Answered by
GPT-4o mini
The atomic mass of an element is a weighted average of the masses of its isotopes, taking into account the abundance of each isotope in nature. Here’s how to calculate the atomic mass:
1. **Identify the Isotopes**: Find the isotopes of the element along with their atomic masses. Atomic masses are usually expressed in atomic mass units (amu).
2. **Determine Abundance**: Obtain the natural abundance (percentage) of each isotope. Abundance values are commonly found in scientific literature or databases.
3. **Convert Abundance**: Convert the percentage abundance of each isotope to a decimal by dividing by 100. For example, if an isotope has an abundance of 75%, its decimal form is 0.75.
4. **Calculate Contributions**: For each isotope, multiply its atomic mass by its decimal abundance. This gives you the contribution of each isotope to the overall atomic mass.
\[
\text{Isotope Contribution} = \text{Atomic Mass of Isotope} \times \text{Decimal Abundance}
\]
5. **Sum Contributions**: Add all the contributions from the isotopes together to get the weighted average atomic mass.
\[
\text{Atomic Mass} = \sum (\text{Isotope Contribution})
\]
### Example:
Let's say an element has two isotopes:
- Isotope 1: Atomic mass = 10 amu, Abundance = 75%
- Isotope 2: Atomic mass = 11 amu, Abundance = 25%
**Calculate Contributions**:
- Isotope 1:
\[
10 \, \text{amu} \times 0.75 = 7.5 \, \text{amu}
\]
- Isotope 2:
\[
11 \, \text{amu} \times 0.25 = 2.75 \, \text{amu}
\]
**Sum Contributions**:
\[
\text{Atomic Mass} = 7.5 \, \text{amu} + 2.75 \, \text{amu} = 10.25 \, \text{amu}
\]
Thus, the atomic mass of the element would be approximately **10.25 amu**.
### Note:
Atomic masses found on the periodic table are typically already calculated and rounded to a certain precision, reflecting the average of naturally occurring isotopes.
1. **Identify the Isotopes**: Find the isotopes of the element along with their atomic masses. Atomic masses are usually expressed in atomic mass units (amu).
2. **Determine Abundance**: Obtain the natural abundance (percentage) of each isotope. Abundance values are commonly found in scientific literature or databases.
3. **Convert Abundance**: Convert the percentage abundance of each isotope to a decimal by dividing by 100. For example, if an isotope has an abundance of 75%, its decimal form is 0.75.
4. **Calculate Contributions**: For each isotope, multiply its atomic mass by its decimal abundance. This gives you the contribution of each isotope to the overall atomic mass.
\[
\text{Isotope Contribution} = \text{Atomic Mass of Isotope} \times \text{Decimal Abundance}
\]
5. **Sum Contributions**: Add all the contributions from the isotopes together to get the weighted average atomic mass.
\[
\text{Atomic Mass} = \sum (\text{Isotope Contribution})
\]
### Example:
Let's say an element has two isotopes:
- Isotope 1: Atomic mass = 10 amu, Abundance = 75%
- Isotope 2: Atomic mass = 11 amu, Abundance = 25%
**Calculate Contributions**:
- Isotope 1:
\[
10 \, \text{amu} \times 0.75 = 7.5 \, \text{amu}
\]
- Isotope 2:
\[
11 \, \text{amu} \times 0.25 = 2.75 \, \text{amu}
\]
**Sum Contributions**:
\[
\text{Atomic Mass} = 7.5 \, \text{amu} + 2.75 \, \text{amu} = 10.25 \, \text{amu}
\]
Thus, the atomic mass of the element would be approximately **10.25 amu**.
### Note:
Atomic masses found on the periodic table are typically already calculated and rounded to a certain precision, reflecting the average of naturally occurring isotopes.