Isotopes of oxygen

Oxygen is a vital element found throughout the universe, but it does not exist in just one form. It comes in various versions called isotopes. Every oxygen atom has eight protons, but the number of neutrons in its nucleus can change. This difference in neutrons creates different isotopes. Some of these are stable, meaning they remain unchanged forever. Others are radioisotopes, which are unstable and undergo radioactive decay. Scientists study these variations to understand everything from the inner workings of stars to the medical imaging of the human brain.

Radioactive isotopes change through specific processes called decay modes. For isotopes lighter than the stable versions, the most common method is positron emission, or $\beta^+$ decay, which turns oxygen into nitrogen. For heavier isotopes, the most common method is beta-minus decay, or $\beta^-$ decay, which turns oxygen into fluorine. Some very heavy isotopes, such as those up to oxygen-28, decay through neutron emission. This process releases a neutron, changing the nucleus into a different element like fluorine-25. These decay paths are predictable and allow scientists to track how atoms transform over time.

There are three known stable isotopes of oxygen: oxygen-16, oxygen-17, and oxygen-18. Oxygen-16 is the most abundant, making up approximately 99.76% of all oxygen. Oxygen-17 is much rarer, with an abundance of about 0.038%. Oxygen-18 is also relatively uncommon, accounting for roughly 0.20% of oxygen. Beyond these stable forms, many radioisotopes exist, ranging from oxygen-11 to oxygen-28. The most stable of these radioisotopes are oxygen-15, with a half-life of 122.27 seconds, and oxygen-14, with a half-life of 70.62 seconds. All other known radioisotopes have even shorter lifetimes.

Many of these isotopes are created through stellar nucleosynthesis, which is the process of making elements inside stars. Oxygen-16 is a major product of stellar evolution. It is synthesized during the helium fusion process. Specifically, the triple-alpha process creates carbon-12, which then captures an additional helium-4 nucleus to form oxygen-16. It can also be created through the neon-burning process. Oxygen-17 is primarily made during the CNO cycle, where stars burn hydrogen into helium. Oxygen-18 is often produced when nitrogen-15 captures a nucleus to become oxygen-16, which then decays into oxygen-18.

Humanity has mapped these isotopes through significant scientific discoveries. Between 1917 and 1919, Frederick Soddy and Ernest Rutherford conducted the first man-made transmutation of nitrogen-14 and helium-4. In 1925, Patrick Blackett hypothesized and imaged oxygen-17. Later, in 1929, Giauque and Johnson detected the presence of oxygen-17 in the Earth's atmosphere using absorption spectra. These milestones proved that oxygen isotopes are not just theoretical, but exist naturally in our world. These discoveries changed how we view the chemical makeup of our planet and the stars.

Oxygen isotopes have immense practical value in medicine and environmental science. Oxygen-15 is used in positron emission tomography, also known as PET scans. Doctors use it for myocardial perfusion imaging of the heart and for brain imaging. Oxygen-18 is vital for paleoclimatology, the study of ancient climates. By measuring the ratio of oxygen-18 to oxygen-16 in ice cores, scientists can calculate historical temperatures. This works because lighter isotopes evaporate more easily than heavier ones. This disparity creates different concentrations in ice and seawater, leaving a record of Earth's past.

Finally, oxygen isotopes connect to broader systems like nuclear energy and biology. In nuclear reactors, the neutron flux can convert oxygen-17 into carbon-14, which is an undesirable product. In biology, oxygen-17 allows for NMR studies of metabolic pathways. This is because it is the only stable oxygen isotope with a nuclear spin. Additionally, lightning can produce oxygen-15 in the air. This happens when gamma rays knock neutrons out of oxygen-16 or nitrogen-14 nuclei. These small, invisible changes help us understand the complex connections between physics, chemistry, and life.