Isotope (${}_Z^A\mathrm{X}$)

Isotopes of a chemical element have the same atomic number $Z$, hence the same number of protons, but different neutron numbers $N$ and therefore different mass numbers $A=Z+N$. For example, every neon atomic nucleus contains ten protons: neon-20, neon-21, and neon-22 contain ten, eleven, and twelve neutrons, respectively. They occupy one chemical element box in the periodic table, rather than three boxes.

Three labels answer different questions. The chemical element fixes $Z$; the isotope specifies which neutron number accompanies that $Z$; the ion label specifies the net electric charge. Neutral ${}_{10}^{22}\mathrm{Ne}$ and the ion ${}_{10}^{22}{\mathrm{Ne}}^{+}$ both contain ten protons and twelve neutrons; they contain ten and nine electrons, respectively. Ordinary ionization changes the electron count, leaving the isotope unchanged. Changing the neutron number can change nuclear binding energy and the possibility of radioactive decay. Isotopic abundances can help distinguish material sources, while radioactive decay can additionally supply chronological information.

A mass spectrum makes the distinction experimentally visible: a single chemical element can give separate ion peaks at different mass-to-charge ratios. Its bulk atomic weight is a weighted mean governed by isotopic abundances, rather than the location of one compulsory peak. The historical mass spectrograph experiment and the synthetic neon mass spectrum below illustrate these complementary observations.

Partitioning of isotopes between materials or processes that changes their relative abundances. Biological and abiotic processes can both produce such differences.

A change in relative abundances of carbon isotopes between materials or processes; both biological and abiotic mechanisms can cause it.

An isotope with no observed radioactive decay on relevant timescales.

The relative abundance of two specified isotopes. Radiogenic growth, mixing and isotope fractionation can modify different ratios in different ways.

An isotopic measurement using the samarium-neodymium parent-daughter system to constrain age or source evolution under specified assumptions.

A ratio of oxygen isotopes, whose differences between minerals and fluids can record temperature and water-rock interaction.

A ratio of hafnium isotopes, often measured in zircon to investigate source extraction and reworking histories.

A ratio of lead isotopes, used with parent-daughter relationships or source mixing models.

A ratio of strontium isotopes; radiogenic ratios can constrain source histories and mixing when initial values and later alteration are considered.

A ratio of boron isotopes, commonly used to study fluid-rock exchange and chemical fractionation.

A radioactive carbon isotope. Its short half-life makes radioactive decay storage relevant to waste-gas management; no biological mechanism is needed for the retention calculation.

Oxygen-18 is the stable oxygen isotope with mass number 18. oxygen from enriched water can trace the source of an oxygen atom in a reaction product, provided subsequent oxygen exchange is excluded.

Measurement of isotopic composition using non-radioactive isotopes. Food-web and dietary interpretations require an appropriate baseline and allowance for tissue formation and physiological fractionation.

A molecular species differing from another only in its isotopic composition. Different reduced masses alter vibrational and rotational spacings even when the underlying electronic potential is nearly the same.

stable isotope labeling uses nonradioactive isotopes to distinguish molecular populations. Paired labeled and unlabeled samples can be compared by mass spectrometry, subject to labeling completeness, mixing and measurement-response controls.

The relative isotopic abundances in a sample.

The mass number $A$ is the total number of nucleons, namely protons plus neutrons, in an atomic nucleus:It is an exact positive integer with no units, rather than a measured mass. Carbon-12 has $A=12$ because its atomic nucleus contains six protons and six neutrons. No improvement in a mass measurement can turn that count into $11.999$ or $12.001$.^[ref]

The left superscript in ${}_{10}^{22}\mathrm{Ne}$ denotes mass number $22$, not an atomic mass of exactly $22\mathrm{u}$. For neon-22, the displayed reference atomic mass is about $21.991385114\mathrm{u}$. Even a weighted mean of mass numbers, such as $20.1877$ for the illustrative neon population below, is a population average, not the mass number of an individual atomic nucleus.^[ref]

The number of neutrons in an atomic nucleus. Two isotopes of the same chemical element differ in neutron number while sharing an atomic number.

The rest mass of a chemical atom in its ground state, a physical quantity expressed in SI mass units or the unified atomic mass unit. Here $m_i$ means the atomic mass of a neutral chemical atom of isotope $i$, not the mass of its bare atomic nucleus and not a population average.^[ref]

Atomic mass includes electrons and depends on nuclear binding energy; it is not just mass number times a universal nucleon mass. The atomic mass of neutral carbon-12 in its ground state is exactly $12\mathrm{u}$ by definition of the unified atomic mass unit. That exactness establishes the reference scale, rather than making the atomic masses of all isotopes exact integers in the same unit.

The unified atomic mass unit and dalton are names for the same mass unit:where the carbon-12 chemical atom is free, at rest, and in its nuclear and electronic ground state. Its conversion to SI mass units is determined experimentally. See The International System of Units (SI), ninth edition, SI Brochure (Table 8, footnote (f), printed p. 146: dalton and unified atomic mass unit).

The mass deficit of a bound atomic nucleus relative to its separated protons and neutrons. If their rest masses are $m_p$ and $m_n$, the atomic nucleus has rest mass $M_{\mathrm{nuc}}$, and the positive nuclear binding energy is $B$, thenwith $c$ the speed of light. Releasing energy when the nucleons bind reduces the rest mass of the bound system without changing the exact count $A=Z+N$. A neutral atomic mass also includes $Z$ electron rest masses, reduced by the energy required to separate those bound electrons. Consequently $A-m_{\mathrm{a}}/\mathrm{u}$ is not the definition of the nuclear mass defect: the reference unit and constituent rest masses must also be accounted for.

For a chemical element in a specified population, its relative atomic mass or atomic weight is the dimensionless ratio $A_{\mathrm{r}}=\bar{m}/\mathrm{u}$, with $\bar{m}$ the mean neutral atomic mass weighted by isotopic abundance. For the illustrative neon population, $A_{\mathrm{r}}\approx 20.1800$, whereas $\bar{m}\approx 20.1800\mathrm{u}$ has mass units.^[ref]

A recommended atomic weight applicable to normal materials, rather than the result for every possible enriched sample. The tabulated neon value $20.1797(6)$ and the arithmetic result above need not agree in every last displayed digit. Some chemical elements have intervals of standard atomic weight because their isotopic abundances vary across normal materials.^[ref]

An isotopic abundance is an atom fraction within a specified population of one chemical element.^[ref] If $n_i$ counts its chemical atoms of isotope $i$, useThe denominator includes every isotope of that chemical element included in the stated population, not every chemical atom of every chemical element in the sample. Thus an oxygen atom fraction among all chemical elements is a different quantity.

For an illustrative neutral neon population, take the displayed central isotopic abundances $f_{20}=0.9048$, $f_{21}=0.0027$, and $f_{22}=0.0925$ from the reference table, so $0.9048+0.0027+0.0925=1$. Its mean atomic mass is the weighted arithmetic meanThese are rounded reference inputs with uncertainties, not exact universal isotopic abundances; the long decimal shows their arithmetic result, not measurement precision.^[ref]

The population convention matters. If instead $w_i$ is the fraction of total neutral-atom mass supplied by isotope $i$, then $w_i=f_i{m}_i/\bar{m}$ andThis is a weighted harmonic mean in the mass fractions. Using ${\sum }_i{w}_i{m}_i$ with those mass fractions would answer a different question. The conversion follows directly by counting chemical atoms and adding their masses.