Which Two Atoms Are Isotopes of One Another? Easy Guide

The quickest way to spot isotopes is to compare two atoms’ nuclear composition. Two atoms are isotopes of one another when they have the same number of protons (the same atomic number, so they are the same element) but different numbers of neutrons, giving them different mass numbers. For example, carbon-12 and carbon-14 each have 6 protons, but 6 and 8 neutrons respectively — so they are isotopes. The proton count fixes the element’s identity; the neutron count only changes the mass.

The Simple Rule for Identifying Isotopes

Every atom is defined by two numbers: its atomic number (protons) and its mass number (protons + neutrons). Isotopes share the first and differ in the second.

So the test is straightforward:

• Same atomic number → same element → possible isotopes.
• Different mass number → different neutron count → they are isotopes.

If two atoms have different proton counts, they are different elements entirely, not isotopes. Hydrogen-1 and helium-4 are not isotopes, because hydrogen has 1 proton and helium has 2.

The term itself comes from the Greek for “same place,” because isotopes of an element occupy the same position on the periodic table even though their masses differ, as noted by the International Union of Pure and Applied Chemistry.

A Worked Example

Compare three atoms:

Carbon-12 and carbon-13 are isotopes: identical protons, different neutrons. Nitrogen-14 has 7 protons, so it is a different element — and not an isotope of either carbon atom, even though its neutron count matches carbon-13’s. Notice that mass number alone tells you nothing about identity; you must check protons first.

Why Neutrons Change but Elements Don’t

Protons carry positive charge and define the element. Neutrons carry no charge; they only add mass and help stabilize the nucleus. Adding or removing a neutron therefore leaves the chemistry essentially unchanged — an isotope of carbon still behaves like carbon in reactions, as Encyclopaedia Britannica explains.

This is why isotopes of an element are chemically almost identical but physically distinct in mass. The differences that matter most are mass, stability, and how easily instruments can separate them.

Stability is a property of the isotope, not a criterion for whether two atoms are isotopes. Carbon-12 and carbon-13 are stable; carbon-14 is radioactive. But carbon-12 and carbon-14 still qualify as isotopes of each other regardless of the fact that one decays.

Common Mistakes to Avoid

People lose marks on this topic in predictable ways. Watch for these traps.

Confusing mass number with atomic number. Mass number is protons plus neutrons; atomic number is protons only. Two atoms with the same mass number but different protons are not isotopes — they are different elements.

Assuming equal neutrons means isotopes. Same neutrons, different protons equals different elements. Isotopes require the protons to match.

Thinking ions are isotopes. An ion is an atom with a gained or lost electron (charge change). Isotopes differ in neutrons (mass change). A quick check: charges involve electrons, isotopes involve neutrons.

Fast Identification Checklist

Run any pair through three questions:

1. Do both atoms have the same number of protons? If no, they are different elements — stop.
2. Do they have different numbers of neutrons (different mass numbers)? If yes, they are isotopes.
3. If protons and neutrons match, they are the same isotope, not two different ones.

Applying this to chlorine-35 and chlorine-37: both have 17 protons, with 18 and 20 neutrons. Same protons, different neutrons — isotopes. Chlorine-35 and argon-36? Protons differ (17 vs. 18), so not isotopes.

How to Read Isotope Notation

Chemists write isotopes in a compact form that encodes both numbers at a glance: the mass number sits as a superscript on the upper left, and the atomic number as a subscript on the lower left. Carbon-14 appears as ¹⁴₆C — the 14 is protons plus neutrons, and the 6 is protons alone. Subtracting gives the neutron count directly: 14 − 6 = 8 neutrons.

This notation makes isotope comparisons fast. If two symbols share the same subscript but have different superscripts, you are looking at isotopes. If the subscripts differ, you are looking at different elements, full stop. Some textbooks drop the subscript entirely because the element symbol already implies the proton count — so ¹⁴C and ¹²C are unambiguously carbon isotopes.

A practical shortcut: the number of neutrons equals mass number minus atomic number. For uranium-235, that is 235 − 92 = 143 neutrons; for uranium-238, it is 238 − 92 = 146. Same 92 protons, three neutrons apart — isotopes.

Isotopes in the Real World

Isotope differences are not just exam material; they drive practical applications across science and medicine.

• Radiometric dating. Carbon-14 forms in the atmosphere and is taken up by living things. When an organism dies, the carbon-14 decays without replacement, so measuring the remaining ratio estimates age.
• Medical imaging. Technetium-99m, a metastable isotope of technetium, is used in nuclear medicine scans because it emits gamma rays that detectors can trace without lingering in the body. The “m” denotes a temporarily excited nuclear state rather than a different proton or neutron count.
• Tracing and forensics. Stable isotope ratios in water, food, and bones vary by geography, letting researchers trace the origins of samples.
• Nuclear energy. Uranium-235 undergoes fission readily, while uranium-238 does not; separating these two isotopes is the central challenge of uranium enrichment, precisely because they are chemically nearly identical.

That last point illustrates why isotopes matter so much: because they behave the same in chemical reactions, separating them requires exploiting the tiny mass difference rather than chemistry.

Hydrogen: The Extreme Case

Hydrogen deserves special mention because its isotopes differ more dramatically than any other element’s. Hydrogen-1 has a single proton and no neutrons, with a mass of about 1 atomic mass unit. Deuterium (hydrogen-2) has one neutron, roughly doubling the mass. Tritium (hydrogen-3) has two neutrons. Because the relative mass jump is so large, deuterium behaves measurably differently from ordinary hydrogen — affecting reaction rates and bond behavior in ways called kinetic isotope effects. This is a striking exception to the “chemically identical” rule of thumb.

Practical Precision

When a question asks which two atoms are isotopes of one another, resist the urge to eyeball the mass numbers. Verify the proton count, confirm the neutron difference, and check that you are not comparing an ion (electron change) or a different element (proton change). Precision matters: a careless reading of the periodic table or a mix-up between atomic number and atomic mass will send you to the wrong answer, especially when multiple answer choices share similar masses.

For health-related applications such as medical imaging, isotope use should always be handled by qualified professionals, and any diagnostic decisions belong with a licensed clinician.

One final habit worth building: always identify the element first, then count neutrons. The element tells you the proton number by definition, so the neutron difference falls out automatically. That two-step routine — element, then neutrons — is the most reliable way to answer any isotope identification question correctly and quickly.

Key Takeaways

• Two atoms are isotopes when they have the same number of protons but different numbers of neutrons.
• Same protons means same element; different mass numbers mean different neutron counts.
• Carbon-12 and carbon-14 are isotopes; carbon-12 and nitrogen-14 are not.
• Mass number alone cannot identify isotopes — always check the proton count first.
• Isotopes share nearly identical chemistry but differ in mass and sometimes in stability (some are radioactive).
• Ions differ by electrons, not neutrons, so an ion is not a different isotope.