Isotope
Atoms of same element differing in neutron number.
Isotopes are distinct nuclear species, or nuclides, of the same chemical element. They share an identical atomic number—the number of protons in the nucleus—and thus occupy the same position in the periodic table, but they differ in their mass number due to varying numbers of neutrons. This means all isotopes of a given element have virtually the same chemical properties, yet they possess different atomic masses and physical characteristics. The term itself derives from the Greek roots for "equal place," reflecting that different isotopes of an element hold the same spot on the periodic table. It was coined by Scottish doctor and writer Margaret Todd in a suggestion to British chemist Frederick Soddy, who popularized the term.
The atomic number, equal to the number of protons, also equals the number of electrons in a neutral atom. While each atomic number defines a specific element, an element's atoms can have a wide range of neutron numbers. The total number of nucleons (protons and neutrons) gives the mass number, which differs for each isotope of an element. For example, carbon has an atomic number of 6, meaning every carbon atom has 6 protons. Its isotopes carbon-12, carbon-13, and carbon-14 have mass numbers 12, 13, and 14, corresponding to neutron numbers of 6, 7, and 8, respectively.
The concept of isotope is closely related to that of a nuclide, which refers to an atom with a specific proton and neutron count. The nuclide concept emphasizes nuclear properties, while the isotope concept groups all atoms of an element and emphasizes chemical properties. The neutron number greatly affects nuclear behavior but has a negligible chemical effect for most elements, though for the lightest element, hydrogen, the isotope effect is large enough to impact biology. Isotopes are specified by the element name followed by a hyphen and the mass number (e.g., helium-3, carbon-14). In standard notation, the mass number appears as a superscript and the atomic number as a subscript to the left of the chemical symbol, though the subscript is often omitted since the element symbol already implies the atomic number. A metastable nuclear isomer is indicated by an appended "m" after the mass number. Some isotopes are radioactive (radioisotopes), while others are stable and have never been observed to decay.
- field
- Nuclear physics, chemistry
- key_concept
- Atoms of same element with different neutron numbers
- popularized_by
- Frederick Soddy
- notable_examples
- Carbon-12, carbon-13, carbon-14
Lore & Background
Isotopes are variants of a chemical element that share the same atomic number—meaning they have the same number of protons in their nuclei and occupy the same position on the periodic table—but differ in their mass number due to a different number of neutrons. For instance, the element carbon always has six protons, but its isotopes carbon-12, carbon-13, and carbon-14 contain six, seven, and eight neutrons respectively. This difference in neutron count gives each isotope a distinct atomic mass and affects physical properties, though chemical behavior remains nearly identical for most elements; the lightest element, hydrogen, is a notable exception where the isotopic effect is large enough to influence biology. The term itself derives from Greek roots meaning "same place," coined by Scottish doctor Margaret Todd in a conversation with British chemist Frederick Soddy, who had observed that radioactive decay could produce chemically identical elements with different masses. Isotopes are specified by the element name followed by a hyphen and mass number (e.g., helium-3, uranium-235) or by standard notation with the mass number as a superscript and atomic number as a subscript to the left of the chemical symbol. Some isotopes are stable, never observed to decay, while others are radioactive and are called radioisotopes; carbon-14, for example, is radioactive, whereas carbon-12 and carbon-13 are stable. Of the roughly 339 naturally occurring nuclides on Earth, 286 are primordial.
Reader's Guide
The concept of isotopes revolutionized understanding of atomic structure and chemical elements. It resolved the puzzle of why different radioactive species could occupy the same place in the periodic table, leading to the recognition that elements could have multiple forms differing only in neutron number. This insight was foundational for nuclear physics, radiochemistry, and later applications in nuclear technology and medicine. The distinction between isotopes and nuclides clarifies that while isotopes group atoms by chemical element, nuclides emphasize nuclear properties. Isotopes have practical importance: radioactive isotopes (radioisotopes) are used in medicine and industry, while stable isotopes help trace chemical and biological processes. The notation system (e.g., carbon-14 or ¹⁴C) standardizes identification. The discovery also explained variations in atomic weights of elements from different sources, as shown by T. W.
Did You Know?
- Carbon-12, carbon-13, and carbon-14 are isotopes of carbon with 6, 7, and 8 neutrons respectively.
- For three elements (tellurium, indium, rhenium) the most abundant natural isotope is actually a long-lived radioisotope.
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