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Atomic Number and Isotopes

Atomic Number and Isotopes

12 min read Updated Mar 26, 2026

These terms sound similar, and the MCAT knows it. Questions that test whether you can distinguish between atomic mass and atomic weight, or between mass number and atomic number, appear regularly. Let’s make the distinctions crystal clear.

Isotope Notation

The standard way to represent an isotope is:

In practice, you will often see isotopes written as the element name followed by the mass number - “carbon-14” or “uranium-238.” Since the element name already tells you the atomic number (carbon is always 6), the mass number is the only additional information you need.

Isotopes

Isotopes are atoms of the same element (same number of protons) that have different numbers of neutrons, and therefore different mass numbers.

Inside an atom, and how to read an isotope symbol

Atomic structure
A schematic atom the electrons occupy all of this volume nucleus p⁺ and n⁰ electrons Not to scale: a real nucleus is ~1/100 000 of the atom's width. The three particles particle charge mass what it decides Proton p⁺ · in the nucleus +1 1.007 amu 1.673 × 10⁻²⁷ kg which element it is Neutron n⁰ · in the nucleus 0 1.009 amu 1.675 × 10⁻²⁷ kg which isotope it is Electron e⁻ · outside, in orbitals −1 0.00055 amu 9.109 × 10⁻³¹ kg the charge, and the chemistry An electron is ~1/1836 of a proton, so an ion weighs what its atom did. Reading the symbol, and the three hydrogens 1 1 H Protium ordinary hydrogen protons (Z) 1 neutrons (A − Z) 0 mass (amu) 1.0078 abundance 99.98 % 2 1 H Deuterium stable, used in D₂O protons (Z) 1 neutrons (A − Z) 1 mass (amu) 2.0141 abundance 0.0156 % 3 1 H Tritium radioactive, t½ ≈ 12.3 y protons (Z) 1 neutrons (A − Z) 2 mass (amu) 3.0160 abundance trace A · mass no. Z · atomic no.
1

Scroll sideways to see the whole map.

Reading the symbolThe bottom number is Z, the proton count, and it is what makes the element what it is. The top number is A, protons plus neutrons. Neutrons are the difference, A − Z, and they are never written directly.

Isotopes versus ionsChanging the neutron count gives an isotope: same element, same chemistry, different mass. Changing the electron count gives an ion: same element, same mass to three figures, different charge and very different behaviour.

Why atomic masses are not whole numbersThe mass on the periodic table is a weighted average over natural abundances. Chlorine sits at 35.45 because it is roughly three quarters chlorine-35 and one quarter chlorine-37, not because any single atom weighs 35.45 amu.

Mass lives in the nucleus, chemistry lives outside it. A proton and a neutron weigh about the same; an electron weighs about one eighteen-hundredth as much, which is why adding or removing electrons changes an atom's charge and reactivity without meaningfully changing its mass.

Consider hydrogen - the simplest element. It has three isotopes:

  • Protium (¹H): 1 proton, 0 neutrons - by far the most abundant
  • Deuterium (²H): 1 proton, 1 neutron - stable, found in “heavy water”
  • Tritium (³H): 1 proton, 2 neutrons - radioactive

All three are hydrogen. All three have one proton. They differ only in neutron count. Because isotopes share the same number of protons and electrons, they have nearly identical chemical properties - they form the same bonds and undergo the same reactions. Their physical properties (mass, radioactive behavior) may differ.

Isotones, Isobars, and Allotropes

The MCAT occasionally tests whether you can distinguish isotopes from three related terms. These show up less frequently than isotopes, but knowing the differences prevents easy points from slipping away.

Isotones are atoms of different elements that have the same number of neutrons. Carbon-14 (6 protons, 8 neutrons) and nitrogen-15 (7 protons, 8 neutrons) are isotones - both have 8 neutrons, but they are entirely different elements.

Isobars are atoms of different elements that have the same mass number. Argon-40 (18 protons, 22 neutrons) and calcium-40 (20 protons, 20 neutrons) are isobars - both have a mass number of 40, but different numbers of protons and neutrons.

Allotropes are different structural forms of the same element in the same physical state. Diamond, graphite, and fullerene are all allotropes of carbon - same element, same number of protons and neutrons, but the atoms are arranged in completely different crystal structures, giving them wildly different physical properties.

| Term | Same Element? | Same Protons? | Same Neutrons? | Same Mass Number? |
|------|:---:|:---:|:---:|:---:|
| Isotopes | Yes | Yes | No | No |
| Isotones | No | No | Yes | No |
| Isobars | No | No | No | Yes |
| Allotropes | Yes | Yes | Yes | Yes |

Atomic Mass vs. Atomic Weight

These two terms are easily confused:

Atomic mass (or mass number, A) is the total count of protons and neutrons in a specific atom. It is always a whole number. Carbon-12 has an atomic mass of 12 amu. Carbon-14 has an atomic mass of 14 amu.

Atomic weight is the weighted average of the atomic masses of all naturally occurring isotopes of an element. It is the number printed on the periodic table, and it is almost never a whole number.

Calculating Atomic Weight

The atomic weight is calculated by multiplying each isotope’s mass by its natural abundance (as a decimal fraction), then summing:

Example: Chlorine has two main isotopes: Cl-35 (75.77%) and Cl-37 (24.23%).

Atomic weight = (35 x 0.7577) + (37 x 0.2423) = 26.52 + 8.97 = 35.49 amu

This is why the periodic table lists chlorine as 35.5, not 35 or 37.

The Mole and Molar Mass

One mole of any substance contains 6.022×10236.022 \times 10^{23} particles (Avogadro’s number, NAN_A). The beauty of the atomic weight is that it connects the microscopic world to the lab bench:

  • The atomic weight of carbon is 12.01 amu
  • One atom of carbon-12 weighs exactly 12 amu
  • One mole of naturally occurring carbon weighs 12.01 grams

In other words, the atomic weight in amu for a single atom equals the molar mass in grams per mole (g/mol) for a mole of atoms. This equivalence is what makes the mole concept so powerful.

Element Q has three isotopes: A (40 amu, 60%), B (44 amu, 25%), and C (41 amu, 15%). What is the atomic weight of Q?
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41.15 amu. (40 x 0.60) + (44 x 0.25) + (41 x 0.15) = 24.00 + 11.00 + 6.15 = 41.15 amu. Notice the result is closest to 40, the most abundant isotope.

What is the difference between atomic mass and atomic weight?
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Atomic mass is the mass of a specific isotope, approximately equal to its mass number (protons + neutrons). Atomic weight is the weighted average of all naturally occurring isotopes, which is the number on the periodic table. Atomic mass is for one isotope; atomic weight is the average across all isotopes.

Carbon-14 (6 protons, 8 neutrons) and nitrogen-15 (7 protons, 8 neutrons) are examples of which relationship: isotopes, isotones, or isobars?
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Isotones. They have the same number of neutrons (8) but different numbers of protons, making them different elements entirely. Isotopes share the same protons (same element). Isobars share the same mass number. Isotones share the same neutron count - the “n” in isotone helps you remember “neutrons.”

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