Isotopes in Astronomy - ChemPRIME

Isotopes in Astronomy

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When the NASA spacecraft Stardust passed alongside Comet Wild 2 in 2006, it used aerogel to capture gas and dust from its coma and tail. When the aerogel was returned to Earth, scientists were excited at the discovery of non-terrestrial glycine, an important amino acid.

However, analysis was required to ensure that the sample was not simply tarnished with Earth glycine en route to the lab. The molecule returned seemed in every way identical to Earth glycine, so NASA analysts were initially puzzled. Space provides a very different chemical environment compared to Earth, but many of these distinctions are quite subtle.

One such distinction is the variation in number of neutrons in atoms of the same element. These species have distinct nuclidic mass, or the mass of a nucleus comprised of protons and neutrons, and are called isotopes. They are usually written symbolically with three components in the MNE format (a simple mnemonic is the word "enemy"). N is the atomic number, number of protons; M stands for the atomic mass, (protons + neutrons); and E is the element symbol. For example, an oxygen atom with 8 protons and 8 neutrons in its nucleus is written as 168O.

Isotopes can occur naturally or artificially through methods such as neutron bombardment. Those rarely seen on Earth play a major role in certain solar reactions, and the ratio of isotopes of a particular element can discriminate between terrestrial and non-terrestrial samples.


Of particular interest to astronomers studying the cosmos for signs of life are extraterrestrial carbon atoms, key components of organic molecules that are the building blocks of life. Because extraterrestrial samples of organic molecules must be returned to Earth to be properly analyzed, and Earth is full of organic molecules, isolating and differentiating a space sample from Earth "contaminants" is key.

The aerogel collector of the Stardust spacecraft.

The identity of carbon atoms, the backbone of all organic chemicals, can help distinguish these samples. On Earth, most C atoms are the familiar Carbon-12, 126C. However, 136C also exists in a 1.1% abundance.[1]. This isotope contains one more neutron than Carbon-12; thus, a 136C atom weighs about 1 amu more than 126C.

Extraterrestrial sources have a slightly higher isotopic ratio of 136C to 126C (about 1.2%)[2]. This seemingly menial percentage difference is substantial when one considers just how many carbon atoms there are in a simple organic molecule, like the glycine collected from Stardust. Fortunately, the isotopic ratio of the glycine sample showed an abnormally high presence of Carbon-13 atoms in the glycine structure, proving that the biomolecule originated in space[3].

In this case, analyzing the mass of the sample was one feasible way to identify between 126C and 136C. A difference in nucleic mass does not necessarily imply a difference in chemical properties. Isotopes have the same number of protons and therefore electrons. Electrons are the main participants in chemical reactions, because they distinguish ions and define chemical bonds. In addition, an atom cannot simply lose or gain an arbitrary number of neutrons, for certain isotopes are very stable and others will decay within a fraction of a second. Measuring the rate of radioactive decay of a sample is another way to determine its isotopic mass. The half-life of the glycine returned from Stardust could have been analyzed and compared to the half-lives of Carbon-12 and Carbon-13.


  2. Mathematically determined from a table at
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