Open any periodic table and you will spot atomic masses like 35.5 or 63.5 — odd, since atoms are built from whole protons and neutrons. The answer is isotopes: atoms of the same element with different masses. For NDA Chemistry this is a small, formula-light topic that delivers guaranteed marks if you learn a handful of definitions and one averaging trick. Let us decode it.
Quick Recap of the Atom
Before isotopes make sense, fix three sub-atomic particles in your head. They are the building blocks examined again and again in NDA.
- Proton — positive charge, found in the nucleus, mass ≈ 1 unit.
- Neutron — no charge (neutral), in the nucleus, mass ≈ 1 unit.
- Electron — negative charge, revolves around the nucleus, mass almost negligible (about 1/1836 of a proton).
Two numbers describe every atom. The atomic number (Z) is the number of protons. The mass number (A) is the total of protons and neutrons.
Atomic number Z = number of protons.
Mass number A = protons + neutrons.
So number of neutrons = A − Z. In a neutral atom, protons = electrons.
An element is defined purely by its atomic number Z. Change Z and you get a different element altogether. Keep Z the same but change the neutrons, and you get an isotope — which is exactly where our story begins.
What Are Isotopes?
Isotopes are atoms of the same element that have the same atomic number (Z) but different mass numbers (A). In plain words, they have the same number of protons but a different number of neutrons.
The word comes from Greek: isos (same) + topos (place). They sit in the same place in the periodic table because they are the same element — only heavier or lighter.
Take the simplest example, hydrogen. It has three isotopes, all with Z = 1 (one proton):
- Protium (1H) — 1 proton, 0 neutrons.
- Deuterium (2H or D) — 1 proton, 1 neutron.
- Tritium (3H or T) — 1 proton, 2 neutrons (radioactive).
Isotopes have identical chemical properties (because chemistry depends on electrons, and the electron count is the same). They differ only in physical properties like mass and density. This is a classic NDA one-liner.
Because they are chemically the same, isotopes of an element cannot be separated by ordinary chemical reactions — only by physical methods that exploit the mass difference, such as diffusion or a mass spectrometer.
It also helps to see why isotopes exist at all. The nucleus is held together by a strong nuclear force acting between protons and neutrons. For a given number of protons, the nucleus can often accommodate a few different neutron counts and still hold together. Each stable neutron count gives a different isotope. Some neutron counts make the nucleus unstable, and those isotopes are radioactive. So an element is really a small family of nuclei sharing the same proton number but carrying slightly different baggage of neutrons.
Common Isotopes You Must Know
NDA loves naming specific isotopes. Memorise this short table of repeat offenders.
- Carbon: 12C, 13C, 14C — all have 6 protons. Carbon-14 is radioactive and used in dating.
- Chlorine: 35Cl and 37Cl — both have 17 protons. This pair explains chlorine’s mass of 35.5.
- Uranium: 235U and 238U — both have 92 protons. U-235 is the fissile fuel for reactors and bombs.
- Oxygen: 16O, 17O, 18O — all have 8 protons.
For any isotope written as AX, instantly read off the neutrons as A − Z. For 37Cl: neutrons = 37 − 17 = 20. Practising this conversion makes most isotope questions a five-second job.
A point students often miss: hydrogen is the only element whose isotopes have separate names — protium, deuterium and tritium. For every other element we simply write the symbol with its mass number, like carbon-12 or uranium-238. Also note that heavy water (D2O), used as a moderator in some nuclear reactors, is just ordinary water in which the hydrogen has been replaced by its heavier isotope deuterium. That single fact links isotopes to the nuclear-reactor questions NDA enjoys asking.
Isotopes vs Isobars vs Isotones
This trio is the single most frequently confused set in the whole chapter. Learn the distinction once and lock it in.
- Isotopes: same atomic number Z (same protons), different mass number A. Example: 35Cl and 37Cl.
- Isobars: same mass number A, different atomic number Z. They are different elements with the same total mass. Example: 40Ar (Z=18) and 40Ca (Z=20).
- Isotones: same number of neutrons, different Z and A. Example: 14C (8 neutrons) and 16O (8 neutrons).
Memory hook:
IsotoPes → same Protons.
IsoBArs → same mass number A (B-A-r).
IsoNes → same Neutrons.
Students assume isobars are isotopes because the masses look equal. They are not — isobars are different elements (different Z). Always check what is kept constant: protons, mass number, or neutrons.
Why Atomic Masses Are Not Whole Numbers
If a single atom of an element has a whole-number mass, why does the periodic table list 35.5 for chlorine?
The answer: most natural elements are a mixture of isotopes. The mass printed in the periodic table is not the mass of one atom — it is the average mass of all the isotopes, weighted by how much of each exists in nature (their relative abundance).
Natural chlorine is roughly 75% chlorine-35 and 25% chlorine-37. Average that and you land close to 35.5, not on any single whole number.
Atomic mass unit (amu / u): defined as exactly 1/12 of the mass of one carbon-12 atom. All relative atomic masses are measured against this standard.
So a fractional atomic mass is actually a clue — it tells you the element exists as a blend of isotopes in fixed proportions. There are exceptions: elements like fluorine, sodium and aluminium exist as a single isotope in nature, so their atomic masses come out very close to whole numbers. The fractional values you usually see are simply the signature of mixed isotopes.
One more idea worth knowing is the difference between atomic mass and atomic weight. In everyday NDA usage they mean the same average value, but strictly the relative atomic mass is dimensionless — it is a comparison against the carbon-12 standard rather than a weight in grams. When you read ‘atomic mass 16 for oxygen’, it means an average oxygen atom is sixteen times heavier than one-twelfth of a carbon-12 atom.
Calculating Average Atomic Mass
This is the only real calculation in the chapter, and it appears regularly. The formula is a simple weighted average.
Average atomic mass = Σ (isotope mass × fractional abundance)
If abundance is given as a percentage, divide by 100 first, or multiply masses by the percentages and then divide the whole sum by 100.
The steps never change:
- Note each isotope’s mass and its % abundance.
- Multiply each mass by its percentage.
- Add the products.
- Divide the total by 100.
If only two isotopes exist and one has abundance x%, the other automatically has (100 − x)%. You are never given both — deduce the second.
Worked Example
Let us pin down chlorine’s famous 35.5 with real numbers.
Chlorine exists as two isotopes: 35Cl with 75% abundance and 37Cl with 25% abundance. Find its average atomic mass.
So chlorine’s atomic mass is 35.5 u — exactly the periodic-table value. No single chlorine atom weighs 35.5; this is the weighted average of the mixture.
An element has two isotopes of masses 10 u and 11 u with abundances 20% and 80%. Find the average atomic mass.
(This is boron, whose listed mass is indeed about 10.8 u.)
Uses of Isotopes
Beyond theory, NDA asks about applications of specific isotopes. Learn these pairings of isotope and use.
- Carbon-14 — radiocarbon dating of fossils, old wood and archaeological remains.
- Cobalt-60 — treatment of cancer (gamma radiation) and sterilising medical equipment.
- Iodine-131 — diagnosing and treating thyroid disorders (goitre).
- Uranium-235 — fuel in nuclear reactors and atomic bombs.
- Phosphorus-32 — tracer to study how plants absorb nutrients and in medical research.
- Sodium-24 — detecting blood clots and circulation problems.
The big three to never forget: C-14 → dating, Co-60 → cancer, I-131 → thyroid. These three have appeared across multiple NDA papers.
Notice that the useful isotopes are mostly radioactive — they emit radiation that we can detect or use to destroy harmful cells. Stable isotopes are mainly studied for their mass differences.
Radioactive Isotopes (Radioisotopes)
Some isotopes have unstable nuclei — too many or too few neutrons to stay together. These break down on their own, releasing radiation. They are called radioactive isotopes or radioisotopes.
The time taken for half of a radioactive sample to decay is its half-life. Carbon-14, for instance, has a half-life of about 5730 years, which is precisely why it dates objects thousands of years old.
Not every isotope is radioactive. 12C and 13C are stable; only 14C is radioactive. So an element can have both stable and unstable isotopes side by side.
Radioactivity was discovered by Henri Becquerel, and the radioactive elements polonium and radium were isolated by Marie and Pierre Curie — names worth keeping handy for general-science questions.
Radioisotopes emit three main kinds of radiation, and NDA sometimes tests these directly. Alpha (α) particles are helium nuclei, heavy and easily stopped by paper. Beta (β) particles are fast electrons with more penetrating power. Gamma (γ) rays are high-energy electromagnetic radiation that can pass through several centimetres of lead. The gamma rays from cobalt-60 are exactly what make it useful for killing cancer cells and sterilising equipment.
Because radioisotopes can be detected even in tiny amounts, they double up as tracers. A small dose of a radioisotope is introduced into a system — a plant, a pipeline, or the human body — and its path is followed using a radiation detector. This is how doctors study blood flow and how engineers locate leaks in underground pipes without digging them up.
Previous-Year Style Practice
Let us apply everything to an exam-style problem.
Q. Atoms of 40Ar (Z = 18) and 40Ca (Z = 20) are an example of which of the following?
Answer: Isobars. Both have the same mass number (A = 40) but different atomic numbers (18 and 20), so they are different elements with equal mass — the very definition of isobars. They are not isotopes, because isotopes must share the same atomic number.
Q. The atomic mass of an element is not always a whole number mainly because of the presence of __________.
Answer: Isotopes. The listed atomic mass is a weighted average of all the isotopes of the element, taken in proportion to their natural abundance, which usually gives a fractional value.
Quick Revision
Run through this list the night before your exam and the chapter is yours.
- Isotopes: same Z (protons), different A. Same chemistry, different mass.
- Isobars: same A, different Z — different elements.
- Isotones: same number of neutrons.
- Neutrons = A − Z; in a neutral atom protons = electrons.
- Atomic mass is a weighted average of isotopes → that is why it is fractional (e.g. Cl = 35.5).
- Average mass = Σ(mass × % abundance) ÷ 100.
- 1 amu = 1/12 of a carbon-12 atom.
- Key uses: C-14 dating, Co-60 cancer, I-131 thyroid, U-235 reactors.
Whenever a question shows two atoms, immediately compare their Z, A and neutron count. That single habit tells you instantly whether you are looking at isotopes, isobars or isotones — and that is most of what NDA asks here.