Enter the atomic mass (in amu) and the relative abundance (as a percentage) for each isotope of the element.
Result:
Note: The total abundance does not equal 100%. Ensure your data is correct.
function calculateAverageAtomicMass() {
var m1 = parseFloat(document.getElementById('mass1').value);
var a1 = parseFloat(document.getElementById('abundance1').value);
var m2 = parseFloat(document.getElementById('mass2').value);
var a2 = parseFloat(document.getElementById('abundance2').value);
var m3 = parseFloat(document.getElementById('mass3').value);
var a3 = parseFloat(document.getElementById('abundance3').value);
var totalMass = 0;
var totalAbundance = 0;
var details = "Calculation: ";
if (!isNaN(m1) && !isNaN(a1)) {
totalMass += (m1 * (a1 / 100));
totalAbundance += a1;
details += "(" + m1 + " × " + (a1 / 100) + ")";
}
if (!isNaN(m2) && !isNaN(a2)) {
totalMass += (m2 * (a2 / 100));
totalAbundance += a2;
details += " + (" + m2 + " × " + (a2 / 100) + ")";
}
if (!isNaN(m3) && !isNaN(a3)) {
totalMass += (m3 * (a3 / 100));
totalAbundance += a3;
details += " + (" + m3 + " × " + (a3 / 100) + ")";
}
if (totalAbundance === 0) {
alert("Please enter valid mass and abundance values for at least two isotopes.");
return;
}
var resultBox = document.getElementById('atomic-result-box');
var massOutput = document.getElementById('mass-output');
var warning = document.getElementById('abundance-warning');
var detailsBox = document.getElementById('calculation-details');
resultBox.style.display = "block";
massOutput.innerHTML = totalMass.toFixed(4) + " amu";
detailsBox.innerHTML = details;
if (Math.abs(totalAbundance – 100) > 0.01) {
warning.style.display = "block";
warning.innerHTML = "Note: Your total abundance is " + totalAbundance.toFixed(2) + "%. Ideally, this should sum to 100%.";
} else {
warning.style.display = "none";
}
}
How to Calculate Average Atomic Mass
The average atomic mass of an element is a weighted average of the masses of its naturally occurring isotopes. Unlike a standard average, a weighted average accounts for how frequently each isotope appears in nature (its relative abundance).
Note: In this formula, the "Abundance" must be expressed as a decimal. To convert a percentage to a decimal, divide it by 100 (e.g., 75% becomes 0.75).
Step-by-Step Calculation Example
Let's look at Chlorine (Cl), which has two main isotopes:
Convert percentages to decimals: 75.78% → 0.7578 and 24.22% → 0.2422.
Multiply mass by decimal abundance:
(34.969 × 0.7578) = 26.499 amu
(36.966 × 0.2422) = 8.953 amu
Sum the results: 26.499 + 8.953 = 35.452 amu.
Why is the Atomic Mass Not a Whole Number?
If you look at the periodic table, you will notice that very few elements have a whole number for their atomic mass. This is because atoms of the same element can have different numbers of neutrons (isotopes). Because these isotopes exist in different proportions in nature, the average mass results in a decimal value. For example, Carbon's atomic mass is 12.011 rather than exactly 12 because of the small percentage of Carbon-13 and Carbon-14 present in nature.
Tips for Accuracy
Significant Figures: Always maintain the number of decimal places provided in your source data until the final step.
Abundance Check: Ensure that the sum of all your fractional abundances equals 1 (or 100%). If it doesn't, you may be missing an isotope or have a typo in your data.
Units: Atomic mass is typically measured in amu (atomic mass units) or u (unified atomic mass units).