Isotopes

Weigh a single hydrogen atom and you will almost always get the same answer — one proton, no neutrons, one electron. But every so often, hiding in the same jar of hydrogen gas, sits an atom that is chemically identical yet noticeably heavier: one proton, but a neutron riding along too. Both atoms are still hydrogen. Both have one proton and one electron, so both react the same way in every test tube experiment you could run. The only thing that differs is the number of neutrons tucked inside the nucleus, and therefore the mass. This is the idea of isotopes: atoms of the same element that share the same atomic number but carry different mass numbers because their neutron counts differ. It explains why the atomic mass printed on your periodic table for chlorine is 35.5 rather than a clean whole number, and it sets up ideas — like carbon-14 dating — that you will meet again later. This article defines isotopes precisely, walks through the standard ICSE examples, and shows how to combine isotopic masses into the average atomic mass you see on the periodic table.

The core insight

Isotopes are atoms of the same element (same atomic number $Z$ , same proton and electron count) that differ in mass number $A$ because they have different numbers of neutrons.

The core idea

Same element, different mass. Every atom of a given element must have the same atomic number $Z$ — that is what makes it that element, since $Z$ fixes the number of protons and (in a neutral atom) the number of electrons. Isotopes are what you get when that same element shows up with different numbers of neutrons packed into the nucleus. The proton and electron counts stay locked; only the neutron count, and hence the total mass, changes. The familiar relationship you already use for any single atom applies separately to each isotope:

A = Z + n

where: $A$ is the mass number (protons plus neutrons), $Z$ is the atomic number (protons, fixed for a given element), and $n$ is the number of neutrons, which is the quantity that varies from one isotope to the next.

Why the chemistry barely changes but the physics does. Chemical behaviour — how readily an atom forms bonds, what compounds it makes, how it reacts — is controlled almost entirely by the arrangement of electrons around the nucleus. Since isotopes of an element have identical electron counts and arrangements, they behave almost identically in chemical reactions. Neutrons, however, sit in the nucleus and add mass without adding charge, so isotopes of the same element can have noticeably different masses and mass-dependent physical properties, such as density or diffusion rate, even though their chemistry is the same.

The three standard families you must know at this level: Hydrogen ( $Z = 1$ ) has three isotopes: protium $^{1}_{1}\text{H}$ (1 proton, 0 neutrons, the ordinary and by far most abundant form), deuterium $^{2}_{1}\text{H}$ (1 proton, 1 neutron), and tritium $^{3}_{1}\text{H}$ (1 proton, 2 neutrons). Carbon ( $Z = 6$ ) has two isotopes named at this level: carbon-12 $^{12}_{6}\text{C}$ (6 protons, 6 neutrons, the reference standard for relative atomic mass) and carbon-14 $^{14}_{6}\text{C}$ (6 protons, 8 neutrons). Chlorine ( $Z = 17$ ) has two isotopes: chlorine-35 $^{35}_{17}\text{Cl}$ (17 protons, 18 neutrons) and chlorine-37 $^{37}_{17}\text{Cl}$ (17 protons, 20 neutrons).

Figure

The three isotopes of hydrogen shown as nuclei with one proton each and zero one or two neutrons, labelled protium deuterium and tritium

Why the periodic table shows fractional masses. No naturally occurring sample of an element is pure one isotope — it is a fixed mixture of its isotopes. The atomic mass printed on the periodic table is a weighted average over that mixture, weighted by how abundant each isotope actually is in nature. That is why chlorine's listed atomic mass, 35.5, is not a whole number: it is not the mass of any single chlorine atom, but an average of chlorine-35 and chlorine-37 atoms in their natural proportions.

Derivation AWeighted average atomic mass from isotopic abundance

The goal is to turn the masses and relative proportions of an element's isotopes into the single average atomic mass value printed on the periodic table, using nothing beyond a weighted average — no calculus, just multiplying each isotope's mass by its share of the sample and adding the results, exactly the arithmetic behind chlorine's familiar value of 35.5.

Step 1. Set up the sample as a fixed group of atoms in the natural ratio. For chlorine, the isotopes chlorine-35 and chlorine-37 occur in roughly a 3-to-1 ratio, meaning that out of every 4 chlorine atoms found in nature, about 3 are chlorine-35 and 1 is chlorine-37. Taking a group of 4 atoms is a convenient way to work with whole numbers before dividing.

Step 2. Add up the total mass contributed by that group. Three chlorine-35 atoms contribute $3 \times 35 = 105$ mass units, and one chlorine-37 atom contributes $1 \times 37 = 37$ mass units, giving a combined total of $105 + 37 = 142$ mass units across the 4 atoms.

Step 3. Divide the total mass by the total number of atoms to get the average mass per atom, which is the weighted average atomic mass:

\text{Average atomic mass} = \frac{(3 \times 35) + (1 \times 37)}{3 + 1} = \frac{142}{4} = 35.5

Step 4. Read the result against the periodic table. The value 35.5 matches the standard listed atomic mass of chlorine exactly, confirming that this weighted-average method is exactly what "atomic mass" means for any element with more than one naturally occurring isotope. The general form of the formula, for an element with isotopes of mass $m_1, m_2, \dots$ present in relative proportions $a_1, a_2, \dots$ , is $\text{Average atomic mass} = \dfrac{\sum (m_i \times a_i)}{\sum a_i}$ , of which the chlorine case is one worked instance.

Worked example

A sample of chlorine gas contains chlorine-35 and chlorine-37 in the approximate natural ratio of 3 atoms of Cl-35 to every 1 atom of Cl-37. Use this ratio to estimate the average atomic mass of chlorine, and state why it is not a whole number.

Step 1. Identify the isotope masses and their ratio: chlorine-35 has mass number 35 and appears 3 parts out of 4, while chlorine-37 has mass number 37 and appears 1 part out of 4.

Step 2. Find the total mass contributed by a representative group of 4 atoms (3 + 1): $(3 \times 35) + (1 \times 37) = 105 + 37 = 142$ .

Step 3. Divide by the total number of atoms in the group to get the weighted average: $\dfrac{142}{4} = 35.5$ .

Step 4. State the reason it is not a whole number: it is an average over two isotopes of different mass, weighted by how common each one is in nature, not the mass of any single chlorine atom. Answer: the average atomic mass of chlorine works out to 35.5, matching the periodic table value, because natural chlorine is a mixture of chlorine-35 and chlorine-37 rather than a single pure isotope.

Common traps

Thinking isotopes are different elements — they are not. Same element means same atomic number; only the neutron count (and hence the mass number) differs between isotopes.
Assuming isotopes react very differently in chemical reactions — their chemistry is nearly identical because chemical behaviour depends on electron number and arrangement, which is the same across all isotopes of an element.
Reading the number in front of an isotope's symbol (like the 14 in carbon-14) as the number of neutrons directly — it is the mass number, and you must subtract the atomic number to get the actual neutron count.
Forgetting that natural samples are mixtures — treating an element's listed atomic mass as the mass of one specific atom, rather than a weighted average over all its naturally occurring isotopes.

Ask the tutor: Want to try more isotope-identification questions, or find out how carbon-14 dating actually works step by step? Ask the AI tutor to walk through a few more examples with you.

If isotopes have the same electrons, why does deuterium behave slightly differently from protium in some experiments?

Isotopes react essentially the same way because chemistry depends on electrons, and this is true for the overwhelming majority of chemical behaviour. But mass does have a very small effect on reaction rates and bond vibration, called an isotope effect, which is why heavy water made with deuterium behaves very slightly differently from ordinary water in some careful measurements. At the ICSE Class 9 level you only need to know that chemical properties are essentially identical; the tiny mass-based differences are advanced content beyond this syllabus.

Are isotopes the same as isobars?

No, and this is a common mix-up. Isotopes are atoms of the same element (same atomic number) with different mass numbers, because their neutron counts differ. Isobars are atoms of different elements that happen to share the same mass number, arrived at by different combinations of protons and neutrons. Isotopes keep the element fixed and vary the neutrons; isobars keep the mass number fixed and vary the element. Isobars are not part of this syllabus, but knowing the distinction prevents a very common exam-answer error.

Why do we care about isotopes like carbon-14 if the syllabus mainly wants definitions?

Carbon-14 is worth knowing about because it becomes useful later: since it is chemically identical to ordinary carbon-12 but radioactively unstable, living things absorb both from their environment, and after death the carbon-14 slowly decays at a known rate. Comparing how much carbon-14 remains in an old sample lets scientists estimate its age, a technique called radiocarbon dating. That application uses radioactivity, which is a separate topic, but it is a good example of why the same-element-different-mass idea genuinely matters outside the classroom.

Isotopes

The core idea

Worked example

Common traps

Before this

Where this leads