Isotopes of oxygen
Oxygen has three stable isotopes: ¹⁶O, ¹⁷O, and ¹⁸O, with ¹⁶O being the most abundant at approximately 99.757% of naturally occurring oxygen, followed by ¹⁸O at 0.205% and ¹⁷O at 0.038%.[1] These isotopes differ in their neutron count—¹⁶O has 8 neutrons, ¹⁷O has 9, and ¹⁸O has 10—while sharing the same 8 protons, leading to variations in atomic mass but similar chemical properties.[1] In addition to these stable forms, oxygen possesses at least 14 known radioactive isotopes, spanning mass numbers from 11 to 28, all of which are highly unstable with half-lives ranging from yoctoseconds to about 2 minutes for the longest-lived, ¹⁵O (half-life 122.2 seconds).[2][1] The stable oxygen isotopes are fundamental in geochemical and environmental studies due to their fractionation during physical processes like evaporation and precipitation, which allows scientists to reconstruct past climate conditions through ratios such as δ¹⁸O (the deviation of ¹⁸O/¹⁶O relative to a standard).[3] For instance, higher δ¹⁸O values in ice cores or sediments indicate warmer temperatures or increased aridity, while lower values suggest colder or wetter periods, enabling paleoclimatologists to trace events like ice ages over millions of years.[3] ¹⁷O also plays a role in advanced triple oxygen isotope analysis (Δ¹⁷O), which provides insights into atmospheric processes, such as the oxygen cycle in air and water, and has applications in studying meteorites and planetary formation.[4] Radioactive oxygen isotopes, though transient, have significant applications in nuclear medicine and research. ¹⁵O is particularly useful in positron emission tomography (PET) scans for measuring cerebral blood flow and oxygen metabolism, as it decays via positron emission to nitrogen-15.[1] Similarly, ¹⁸O serves as a precursor for producing fluorine-18, a key radionuclide in fluorodeoxyglucose (FDG)-PET imaging for cancer detection.[1] Other short-lived isotopes like ¹⁴O and ¹⁹O are employed in nuclear physics experiments to probe neutron-rich nuclei and reaction mechanisms, contributing to our understanding of stellar nucleosynthesis.[5][6] Overall, the isotopes of oxygen underpin diverse fields from Earth sciences to medicine, leveraging both their natural abundances for tracing environmental changes and their radioactivity for diagnostic and experimental purposes.[7]Overview
Basic properties
Oxygen isotopes are nuclides of the element oxygen, characterized by 8 protons in the nucleus and varying numbers of neutrons from 3 to 17, resulting in mass numbers ranging from ¹¹O to ²⁵O. These isotopes exhibit diverse nuclear structures, with stability determined by the balance between protons and neutrons, as well as shell effects in light nuclei. The known oxygen isotopes span this range, though the lightest (¹¹O–¹²O) and heaviest (²³O–²⁵O) are unbound or highly unstable, decaying rapidly via particle emission or beta decay.[8] The nuclear binding energy of oxygen isotopes, which holds the nucleus together against the repulsive Coulomb forces between protons, follows trends well-approximated by the semi-empirical mass formula (SEMF). The SEMF expresses the binding energy B(A, Z=8) as B(A, 8) = a_v A - a_s A^{2/3} - a_c \frac{8(8-1)}{A^{1/3}} - a_a \frac{(A - 16)^2}{A} \pm \delta, where a_v \approx 15.5 MeV is the volume term (favoring larger A), a_s \approx 16.8 MeV the surface term (penalizing small A), a_c \approx 0.72 MeV the Coulomb term (increasing with fixed Z and larger A), a_a \approx 23 MeV the asymmetry term (penalizing deviation from N=Z), and \delta the pairing term (positive for even-even nuclei like ¹⁶O and ¹⁸O, negative for odd-odd, zero for odd A). For oxygen isotopes, the binding energy per nucleon B/A increases from low values in proton-rich isotopes (e.g., ~5–6 MeV for ¹²O) to a peak of 7.976 MeV at ¹⁶O, then decreases to ~7.6 MeV at ²⁰O due to growing asymmetry and surface effects; this trend underscores why only ¹⁶O, ¹⁷O, and ¹⁸O are stable.[9] Spin and parity values for oxygen isotopes reflect the shell model structure, with even-even nuclei (even N, even Z) typically having ground-state spin-parity 0⁺ due to pairing, while odd-N isotopes like ¹⁷O (5/2⁺) and ¹⁹O (5/2⁺) arise from an unpaired neutron in the 1p_{1/2} or similar orbital. Examples among unstable isotopes include ¹⁵O (1/2⁻, from 1s_{1/2} hole) and ²⁰O (0⁺, even-even but weakly bound). These assignments provide introductory insights into nuclear stability, with even-even configurations generally more bound.[8] The following table summarizes key properties for all known oxygen isotopes (¹¹O to ²⁵O), including mass number A, neutron number N, ground-state spin-parity J^π, and stability status (stable or unstable; natural abundances for stable isotopes are 99.757% for ¹⁶O, 0.038% for ¹⁷O, and 0.205% for ¹⁸O). Basic decay thresholds indicate whether the isotope is bound (stable) or unbound/decays promptly (unstable), without detailed modes.[8][2][10][11]| Mass Number (A) | Neutron Number (N) | Spin and Parity (J^π) | Stability Status |
|---|---|---|---|
| 11 | 3 | (3/2⁻) | Unstable |
| 12 | 4 | 0⁺ | Unstable |
| 13 | 5 | (3/2⁻) | Unstable |
| 14 | 6 | 0⁺ | Unstable |
| 15 | 7 | 1/2⁻ | Unstable |
| 16 | 8 | 0⁺ | Stable (99.757%) |
| 17 | 9 | 5/2⁺ | Stable (0.038%) |
| 18 | 10 | 0⁺ | Stable (0.205%) |
| 19 | 11 | 5/2⁺ | Unstable |
| 20 | 12 | 0⁺ | Unstable |
| 21 | 13 | (5/2⁺) | Unstable |
| 22 | 14 | 0⁺ | Unstable |
| 23 | 15 | 1/2⁺ | Unstable |
| 24 | 16 | 0⁺ | Unstable |
| 25 | 17 | 3/2⁺ | Unstable (unbound) |
Natural occurrence
Oxygen has three stable isotopes in nature: ^{16}O, ^{17}O, and ^{18}O, with primordial abundances in Earth's atmosphere and standard mean ocean water (SMOW) of 99.757% for ^{16}O, 0.038% for ^{17}O, and 0.205% for ^{18}O.[12] These values represent the baseline isotopic composition before significant fractionation occurs, as defined by international standards for terrestrial materials.[12] Isotopic fractionation of oxygen arises from both equilibrium and kinetic processes, leading to variations in the ratios of heavier to lighter isotopes. Equilibrium fractionation is temperature-dependent and occurs during reversible reactions, such as mineral-water exchanges, where heavier isotopes like ^{18}O preferentially bond in solids or liquids over ^{16}O.[13] Kinetic fractionation, in contrast, results from irreversible processes like evaporation and diffusion, where lighter isotopes move faster, depleting heavy isotopes in the remaining phase; for example, water vapor becomes enriched in ^{16}O during evaporation, while precipitation is relatively depleted.[13] These mechanisms alter the δ^{18}O and δ^{17}O values, defined relative to SMOW, with δ^{18}O often serving as a proxy for temperature and δ^{17}O revealing additional fractionation details due to its mass-dependent behavior.[14] In natural reservoirs, oxygen isotope ratios vary across the atmosphere, hydrosphere, and biosphere. Atmospheric O_2 exhibits a δ^{18}O value about 23.5‰ higher than mean ocean water due to the "Dole effect," driven by photosynthetic discrimination against heavier isotopes in plants and subsequent respiration.[15] The water cycle introduces spatial variations, with evaporated water vapor depleted in ^{18}O by up to 10-20‰ compared to source water, while polar ice cores show even greater depletions (δ^{18}O as low as -50‰) reflecting colder temperatures and Rayleigh distillation.[13] Biospheric influences, including photosynthesis and microbial activity, further modify ratios; for instance, leaf water enrichment in ^{18}O by 10-30‰ affects atmospheric O_2 through oxygen production.[15] Anthropogenic activities, particularly fossil fuel burning, indirectly impact oxygen isotope ratios by altering atmospheric CO_2 and O_2 dynamics. Combustion consumes O_2 in a stoichiometric ratio of about 1.42:1 relative to CO_2 produced, slightly depleting atmospheric O_2 and influencing its δ^{18}O through feedbacks on global respiration and photosynthesis rates.[15] The oxygen in fossil-derived CO_2, often depleted in ^{18}O compared to natural sources, dilutes the atmospheric CO_2 δ^{18}O by approximately 0.3-0.5‰ since pre-industrial times, as CO_2 exchanges with ocean and terrestrial waters without direct O_2 involvement.[15]Stable isotopes
Oxygen-16
Oxygen-16 (¹⁶O) is the most abundant and lightest stable isotope of oxygen, comprising approximately 99.76% of all naturally occurring oxygen atoms.[16] Its precise atomic mass is 15.994914619(1) u, which serves as a foundational value in atomic weight standards.[17] This isotope's low mass contributes significantly to the physical properties of oxygen-containing compounds, particularly the density of water, where Standard Mean Ocean Water (SMOW)—dominated by ¹⁶O—exhibits a density of 999.975 kg/m³ at 4°C and 1 atm.[18] Variations in heavier oxygen isotopes relative to ¹⁶O can alter water density by up to 0.1126 g/cm³ per unit increase in the mole fraction of ¹⁸O, highlighting ¹⁶O's role as the baseline for such measurements.[18] As the reference isotope in stable oxygen isotope geochemistry, ¹⁶O forms the denominator in ratio measurements for δ¹⁸O and δ¹⁷O notations.[19] These are defined by the formula \delta = \left( \frac{R_{\text{sample}}}{R_{\text{standard}}} - 1 \right) \times 1000‰, where R = {}^{18}\text{O}/^{16}\text{O} or {}^{17}\text{O}/^{16}\text{O}, and the standard is Vienna Standard Mean Ocean Water (VSMOW), which is calibrated against ¹⁶O-rich compositions.[20] This standardization enables precise comparisons of isotopic compositions across samples, with ¹⁶O providing the invariant light reference.[19] ¹⁶O predominates in biological and geological systems, forming the majority of oxygen atoms in essential molecules such as water (H₂O) and carbon dioxide (CO₂), as well as in rock-forming silicates like quartz and feldspar.[16] Its near-universal presence—over 99.7% in terrestrial oxygen reservoirs—ensures that these systems reflect ¹⁶O's properties in bulk analyses.[17] Due to its lowest mass among stable oxygen isotopes, ¹⁶O undergoes minimal equilibrium fractionation in many physicochemical processes, often serving as the "light" end-member in isotope ratios where heavier isotopes are preferentially enriched in condensed phases.[21] This characteristic minimizes deviations in ¹⁶O abundance during evaporation or precipitation, stabilizing its role in ratio-based studies. Compared to the trace abundances of ¹⁷O (0.038%) and ¹⁸O (0.20%), ¹⁶O's dominance also underpins the detection of mass-independent fractionation anomalies in atmospheric and extraterrestrial samples.[16]Oxygen-17
Oxygen-17 (¹⁷O) is a stable isotope of oxygen with an atomic mass of 16.99913175650(69) u and a natural abundance of approximately 0.038%.[12] This low abundance makes it the rarest of the three stable oxygen isotopes, which co-occur in water and other natural samples alongside ¹⁶O and ¹⁸O, often analyzed via mass spectrometry.[7] The nucleus of ¹⁷O has a spin quantum number I = 5/2, which imparts magnetic properties that enable its detection through nuclear magnetic resonance (NMR) spectroscopy.[22] This quadrupolar nucleus allows researchers to probe the chemical environments of oxygen atoms in molecules, providing insights into bonding and structure in both organic and inorganic compounds, despite challenges from low sensitivity and broad linewidths due to the quadrupole moment.[23] In triple oxygen isotope systematics, ¹⁷O plays a key role through the parameter Δ¹⁷O, defined as Δ¹⁷O = δ¹⁷O - 0.52 × δ¹⁸O, which quantifies mass-independent fractionation effects.[24] This metric is essential in atmospheric chemistry for tracing ozone-related processes and in meteoritics for identifying presolar materials and planetary formation histories.[25] Due to its scarcity, samples enriched in ¹⁷O are prepared using techniques such as fractional distillation of water or chemical exchange reactions to achieve higher concentrations for research purposes.[26][27] These methods exploit differences in physical or chemical properties to separate isotopes, enabling detailed studies that would be impractical with natural abundance levels.[28]Oxygen-18
Oxygen-18 (¹⁸O) is the heaviest stable isotope of oxygen, with an atomic mass of 17.9991610 u and a natural abundance of 0.205% in Earth's atmosphere and hydrosphere.[12] It consists of 8 protons and 10 neutrons, making it two neutrons heavier than the dominant isotope, oxygen-16 (¹⁶O), to which ¹⁸O ratios serve as the international standard (VSMOW) for stable isotope measurements.[12] Unlike radioactive isotopes, ¹⁸O does not decay and persists indefinitely, enabling its use as a tracer in geological and biological processes. Oxygen-18 is also analyzed alongside oxygen-17 in triple oxygen isotope studies to distinguish mass-dependent from mass-independent fractionations. The fractionation of ¹⁸O relative to ¹⁶O arises primarily from differences in zero-point energies (ZPE) of molecular bonds, leading to preferential incorporation of the heavier isotope into condensed phases at lower temperatures. The equilibrium fractionation factor \alpha is approximated by \alpha = \frac{k_{\text{heavy}}}{k_{\text{light}}} \approx \exp\left(\frac{\Delta E}{RT}\right), where \Delta E represents the ZPE difference between isotopologues, R is the gas constant, and T is temperature in Kelvin; this relationship underpins temperature-dependent proxies in paleoclimatology. For instance, during evaporation, lighter ¹⁶O preferentially enters the vapor phase, enriching residual water in ¹⁸O, while condensation reverses this, with heavier isotopes condensing first. This behavior is quantified via \delta^{18}\mathrm{O} values (in ‰ relative to VSMOW), where positive values indicate ¹⁸O enrichment and negative values depletion. Variations in \delta^{18}\mathrm{O} serve as key environmental proxies for reconstructing past climates, reflecting temperature, precipitation sources, and ice volume changes. In Antarctic ice cores like Vostok, \delta^{18}\mathrm{O} shifts by approximately 5‰ between glacial and interglacial periods, with more depleted values (e.g., -60‰) during colder glacials due to increased fractionation in colder source regions and Rayleigh distillation effects. Similarly, in marine sediments, benthic and planktonic foraminifera record \delta^{18}\mathrm{O} variations of 2–3‰ across glacial-interglacial cycles, combining signals from ocean temperature (∼1.5‰ per 10°C) and global ice volume (∼1.2‰ for full glaciation).[29] Tree-ring cellulose \delta^{18}\mathrm{O} integrates seasonal precipitation and temperature, showing correlations with relative humidity and source water composition over centuries, as seen in mid-latitude chronologies spanning the last millennium.[30] In biosynthetic pathways, ¹⁸O enrichment occurs differentially during carbohydrate metabolism, with sugars exhibiting partial oxygen exchange with cellular water (∼40–50% of atoms), leading to \delta^{18}\mathrm{O} values intermediate between source water and fully exchanged products. Cellulose, synthesized from these sugars, shows further enrichment (typically +27‰ relative to source water) due to complete exchange at synthesis sites and kinetic fractionation in enzymatic steps, contrasting with less exchanged lipids or proteins.[31] This biochemical fractionation preserves environmental signals in plant archives, aiding reconstructions of hydrological cycles.Radioactive isotopes
Light isotopes (A ≤ 15)
The light isotopes of oxygen with mass numbers A ≤ 15 are proton-rich nuclides situated near or beyond the proton drip line, exhibiting extreme instability due to insufficient binding energy for the excess protons, resulting in half-lives spanning from femtoseconds to minutes. These isotopes decay predominantly by positron emission (β⁺) or electron capture (EC), with the lightest undergoing proton emission or two-proton (2p) decay as they are unbound ground states or low-lying resonances. Their study is crucial for understanding nuclear shell structure in the p-shell region, where Coulomb repulsion dominates over the strong force, leading to broad decay widths and challenges in theoretical modeling such as ab initio calculations or shell model approaches. The known isotopes in this range include ^9O to ^15O, though ^9O, ^10O, and ^11O are unbound and decay almost instantaneously by proton emission for ^9O and ^11O or two-proton emission for ^10O with lifetimes on the order of 10^{-21} to 10^{-15} s, providing limited experimental data beyond resonance parameters. ^12O, also unbound, decays by 2p emission to ^10C with a resonance lifetime of approximately 1 ps, as determined from fragmentation experiments, illustrating the transition from sequential to simultaneous two-proton emission in light nuclei. ^13O, with a half-life of 8.58 ms, undergoes β⁺ decay to ^13N, populating excited states in the daughter. ^14O has a half-life of 70.62 s and decays by β⁺ emission to ^14N, with branching ratios favoring the ground state transition consistent with conserved vector current hypothesis. ^15O is the longest-lived, with a half-life of 122.24 s, decaying 99.9% by β⁺ to ^15N (Q_β = 1.732 MeV) and 0.1% by EC.[32][33][34]| Isotope | Half-life | Primary decay mode | Daughter nuclide |
|---|---|---|---|
| ^9O | ~20 fs (resonance) | p | ^8N |
| ^10O | ~10 fs | 2p | ^8C |
| ^11O | ~0.1 ps | p | ^10N |
| ^12O | ~1 ps | 2p | ^10C |
| ^13O | 8.58 ms | β⁺ | ^13N |
| ^14O | 70.62 s | β⁺ | ^14N |
| ^15O | 122.24 s | β⁺ (99.9%) | ^15N |
Intermediate isotopes (A = 16–19)
The only radioactive ground-state isotope in this mass range is ^{19}O, characterized by a relatively long half-life of 26.470 ± 0.018 seconds compared to lighter oxygen radioisotopes. This half-life allows for laboratory studies of its decay properties, distinguishing it from the shorter-lived species in lower mass regions that exhibit more exotic decay modes like β⁺ emission or proton drip-line behavior.[36][37] ^{19}O undergoes β⁻ decay to ^{19}F with a Q-value of 4.821 MeV, primarily populating excited states in the daughter nucleus rather than the ground state. The dominant branches feed the 197 keV (5/2⁺) state at approximately 30–45% and the 1554 keV (3/2⁺) state at 54–70%, while the branch to the ground state (1/2⁺) is limited to ≤4%; smaller contributions (0.05–7%) occur to other low-lying levels such as the 110 keV (1/2⁻) and 1346 keV (5/2⁻) states. These log ft values, ranging from 4.33 for the strongest Gamow-Teller transition to >6.5 for forbidden branches, reflect a mix of allowed and forbidden decay modes transitional between the proton-rich and neutron-rich extremes. De-excitation in ^{19}F proceeds via γ emission, with key lines at 197 keV and 1356 keV observed in experiments.[36][38] In laboratory production, ^{19}O is generated via charged-particle reactions on stable oxygen targets, notably the ^{18}O(d,p)^{19}O transfer reaction using deuteron beams on enriched ^{18}O, which populates the ground state with high selectivity for decay studies. This method, explored in accelerator experiments, enables precise measurement of spectroscopic factors and asymptotic normalization coefficients relevant to indirect capture rates. In astrophysical contexts, ^{19}O plays a role in inhomogeneous Big Bang nucleosynthesis models, where the ^{18}O(n,γ)^{19}O neutron capture rate influences light element yields beyond standard homogeneous predictions; its subsequent β⁻ decay contributes to fluorine production in primordial scenarios.[39]Heavy isotopes (A ≥ 20)
The heavy isotopes of oxygen, spanning mass numbers from 20 to 25, represent neutron-rich nuclides far from stability, where the addition of excess neutrons leads to rapid β⁻ decay and, near the neutron drip line, neutron emission or unbound resonances. These isotopes provide key insights into nuclear structure at extreme neutron-to-proton ratios, particularly the evolution of shell closures in the sd-shell region. Production of these short-lived species occurs primarily through high-energy reactions, such as multinucleon transfer or fragmentation in heavy-ion collisions at relativistic energies, as demonstrated in experiments at facilities like GSI and NSCL. In astrophysical contexts, they contribute to nucleosynthesis pathways in core-collapse supernovae, where rapid neutron capture (r-process) conditions can transiently form such neutron-excessive light nuclei before further processing.[40][41] Half-lives decrease markedly with increasing neutron number, reflecting heightened instability from the neutron excess that weakens binding energies. For instance, ^{20}O undergoes β⁻ decay to ^{20}F with a half-life of 13.51(5) s, while ^{21}O similarly decays by β⁻ emission in 3.42(10) s. The trend continues with ^{22}O, which has a half-life of 2.25(9) s and decays predominantly (>78%) by β⁻ to ^{22}F, with a minor branch (<22%) for β⁻ delayed neutron emission to ^{21}F. Lighter isotopes like those with A ≤ 15 are proton-rich and favor β⁺ or proton decay, contrasting the β⁻ dominance here due to neutron richness.[32][32][32] Further along the chain, stability erodes rapidly: ^{23}O decays by β⁻ with a half-life of 82(2) ms, and ^{24}O, the heaviest bound oxygen isotope, has a half-life of 77(4) ms, primarily via β⁻ to ^{24}F, though studies of its decay reveal sequential two-neutron emission channels in correlated three-body final states (^{24}O → ^{22}O + 2n). Beyond this, ^{25}O is unbound, manifesting as a low-lying resonance in the ^{24}O + n continuum with resonance energy 0.75(8) MeV and width 10(8) keV above the drip line, populated via proton knockout from ^{26}F; its effective "lifetime" is on the order of picoseconds due to immediate neutron decay. These properties highlight the neutron drip-line boundary, where one- or two-neutron separation energies approach zero.[32][32][42] Shell model interpretations reveal structural nuances, notably a subshell closure at N=14 in ^{22}O (Z=8), where the 2₁⁺ excited state at 3.21 MeV indicates enhanced stability relative to neighbors, forming part of an "island of inversion" boundary where deformation intrudes on shell-model predictions. This closure arises from the filling of the 0d_{5/2} neutron orbital, influencing binding and excitation spectra across the oxygen chain. Observations of these isotopes in cosmic rays further probe high-energy astrophysical acceleration and fragmentation processes.[40]| Isotope | Half-life | Primary Decay Mode | Daughter Product |
|---|---|---|---|
| ^{20}O | 13.51(5) s | β⁻ | ^{20}F |
| ^{21}O | 3.42(10) s | β⁻ | ^{21}F |
| ^{22}O | 2.25(9) s | β⁻ (78%), β⁻n (22%) | ^{22}F, ^{21}F |
| ^{23}O | 82(2) ms | β⁻ | ^{23}F |
| ^{24}O | 77(4) ms | β⁻ | ^{24}F |
| ^{25}O | Unbound resonance (Γ ≈ 10 keV) | n emission | ^{24}O + n |