Atomic Mass: The Weight of the Microscopic World
The Building Blocks of Matter
Everything you see around you — the air you breathe, the water you drink, the device you're reading this on — is made of atoms. Atoms are the basic units of matter. For a long time, people thought atoms were the smallest possible particles, but we now know they are made up of even smaller particles called subatomic particles. The three main ones are:
- Protons: Positively charged particles found in the atom's center, called the nucleus.
- Neutrons: Neutral particles (no charge) that also reside in the nucleus.
- Electrons: Negatively charged particles that orbit the nucleus at high speeds.
The nucleus, containing protons and neutrons, is the dense, heavy core of the atom. The electrons are incredibly light and move in a large space around the nucleus. Because almost all of an atom's mass is concentrated in the nucleus, the atomic mass is essentially the mass of the protons plus the mass of the neutrons.
$ \text{Atomic Mass} \approx \text{(Number of Protons)} + \text{(Number of Neutrons)} $
This sum is also known as the Mass Number.
What is an Atomic Mass Unit (amu)?
Atoms are far too small to be weighed on a regular scale. If we used grams or kilograms, the numbers would be incredibly tiny and hard to work with. For example, a single proton has a mass of about 0.00000000000000000000000167 grams! To make things easier, scientists created a special unit called the atomic mass unit (amu) or unified atomic mass unit (u).
One atomic mass unit is defined as exactly one-twelfth the mass of a carbon-12 atom. A carbon-12 atom has 6 protons, 6 neutrons, and 6 electrons. By international agreement, this atom is assigned a mass of exactly 12 amu.
This means:
- 1 proton has a mass of approximately 1.007 amu.
- 1 neutron has a mass of approximately 1.009 amu.
- 1 electron has a mass of only about 0.00055 amu, which is so small that we usually ignore it when calculating atomic mass.
Mass Number vs. Atomic Mass
It's easy to confuse mass number with atomic mass, but they are different concepts.
Mass Number is a simple count of the protons and neutrons in a specific atom's nucleus. It is always a whole number. For any element, the number of protons (the atomic number) is fixed, but the number of neutrons can vary.
Atomic Mass (also called atomic weight) is the average mass of all the naturally occurring atoms of an element, measured in atomic mass units (amu). Because it is an average, it is usually not a whole number.
Let's look at an example with Carbon. Carbon always has 6 protons. Its atomic number is 6. But it can have different numbers of neutrons.
| Isotope Name | Protons | Neutrons | Mass Number | Atomic Mass (amu) |
|---|---|---|---|---|
| Carbon-12 | 6 | 6 | 12 | 12.000 (exact) |
| Carbon-13 | 6 | 7 | 13 | 13.003 |
The Role of Isotopes
Atoms of the same element that have different numbers of neutrons are called isotopes1. As shown in the table above, Carbon-12 and Carbon-13 are isotopes of carbon. Most elements have several naturally occurring isotopes.
The atomic mass listed on the periodic table is a weighted average of the masses of all the naturally occurring isotopes. The "weighted" part is important because some isotopes are much more common than others.
For example, about 98.93% of all carbon atoms are Carbon-12, and about 1.07% are Carbon-13. To find the average atomic mass of carbon, we do a weighted average calculation:
$ (0.9893 \times 12.000 \text{ amu}) + (0.0107 \times 13.003 \text{ amu}) = 12.011 \text{ amu} $
This is why the atomic mass of carbon on the periodic table is 12.011 amu, not 12.000.
Calculating Atomic Mass: A Step-by-Step Example
Let's take the element Chlorine (Cl) as a practical example. Chlorine has two major isotopes:
- Chlorine-35: Mass = 34.969 amu, and its natural abundance is 75.78% (0.7578).
- Chlorine-37: Mass = 36.966 amu, and its natural abundance is 24.22% (0.2422).
We calculate the average atomic mass as follows:
Step 1: Multiply the mass of each isotope by its abundance (as a decimal).
For Chlorine-35: $ 34.969 \times 0.7578 = 26.50 $
For Chlorine-37: $ 36.966 \times 0.2422 = 8.95 $
Step 2: Add the results together.
$ 26.50 + 8.95 = 35.45 $
So, the atomic mass of chlorine is 35.45 amu, which matches the value on the periodic table. Notice how the value is much closer to 35 because the lighter isotope is more common.
Common Mistakes and Important Questions
Q: Is atomic mass the same as mass number?
A: No. Mass number is a simple count of protons and neutrons in a specific isotope, and it is always a whole number. Atomic mass is the weighted average mass of all naturally occurring isotopes of an element, measured in amu, and it is almost never a whole number.
Q: Why do we ignore the mass of electrons when calculating atomic mass?
A: A proton and a neutron each have a mass that is about 1,836 times greater than the mass of an electron. Because electrons are so incredibly light, their total contribution to the atom's mass is negligible, much like a single feather's weight on a large truck. For all practical purposes, the mass of an atom comes from its nucleus.
Q: If a proton and a neutron each weigh about 1 amu, why isn't the atomic mass always a whole number?
A: There are two main reasons. First, as we saw with the isotopes, the atomic mass is an average of different masses. Second, there is a phenomenon called the mass defect2. The actual mass of a nucleus is always slightly less than the sum of the masses of its individual protons and neutrons. This "lost" mass is converted into the energy that holds the nucleus together, as described by Einstein's famous equation, $E=mc^2$.
Footnote
1 Isotopes: Variants of a particular chemical element which have the same number of protons but different numbers of neutrons.
2 Mass Defect: The difference between the mass of an atom and the sum of the masses of its individual protons, neutrons, and electrons. This "missing" mass is converted into binding energy that holds the nucleus together.