In this case, one of the neutrons in carbon changes into a proton, forming nitrogen During this process, which is known as beta decay , the nucleus emits radiation in the form of an electron and an antineutrino.
There are many factors that can cause a nucleus to decay. One of the most important is the ratio of protons to neutrons a particular nucleus has. The same is true if a nucleus has too many protons. This is one of the reasons why some isotopes of a given element are radioactive, while others are not. By now, you may be wondering how all these isotopes were created in the first place.
As it turns out, this question is a complex one, but lends some truth to the adage that we are all made of star dust. Some of the lighter isotopes were formed very early in the history of the universe, during the Big Bang.
Others result from processes that happen within stars or as a result of chance collisions between highly energetic nuclei - known as cosmic rays - within our atmosphere.
Most naturally existing isotopes are the final stable or long-lived product resulting from a long series of nuclear reactions and decays. In most of these cases, light nuclei have had to smash together with enough energy to allow the strong force - a glue-like bond that forms when protons and neutrons get close enough to touch - to overcome the electromagnetic force — which pushes protons apart. If the strong force wins out, the colliding nuclei bind together, or fuse, to form a heavier nucleus.
Our sun is a good example of this. One of its main sources of power is a series of fusion reactions and beta decay processes that transform hydrogen into helium. Since the early s, when the existence of isotopes was first realised, nuclear physicists and chemists have been seeking out ways to study how isotopes can be formed, how they decay, and how we might use them.
As it turns out, the nature of isotopes — their chemical uniformity, their nuclear distinctiveness — makes them useful for a wide range of applications in fields as diverse as medicine, archaeology, agriculture, power generation and mining.
To find out how many neutrons an isotope harbors, subtract its atomic number from its mass number. For the most part, no. Which is just peachy for us. Taken together, the 81 stable elements known to us can boast some stable isotopes.
Imagine the headache it would cause if they all behaved in a different way. Carbon itself has 3 stable isotopes — would we even exist today if each had its own quirks? One element whose isotopes do differ meaningfully, however, is the runt of the periodic table: hydrogen.
Hydrogen is the simplest chemical element, one proton orbited by one electron. In absolute terms, though, the difference is immense: one neutron will double the mass of a hydrogen atom — two neutrons will triple it.
For comparison, a single neutron is just While isotopes are highly similar chemically, they do differ physically. All that weight can alter how isotopes of light elements, hydrogen especially, behave. One example of such differences is the kinetic isotope effect — basically, heavier isotopes of the same element tend to be more sluggish during chemical reactions than lighter isotopes.
For heavier elements, this effect is negligible. So, molecules that contain isotopes will look different to the same molecule sans isotopes when seen through an infrared camera. This, agian, is caused by their extra mass — the shape and masses of atoms in a molecule change how it vibrates, which in turn, changes how they interact with photons in the infrared range.
Long story short, isotopes are simply atoms with more neutrons — they were either formed that way, enriched with neutrons sometime during their life, or are originated from nuclear processes that alter atomic nuclei. So, they form like all other atoms. Lighter isotopes likely came together a bit after the Big Bang, while heavier ones were synthesized in the cores of stars.
Isotopes can also form following the interaction between cosmic rays and energetic nuclei in the top layers of the atmosphere. These atoms are called radioactive isotopes or radioisotopes. Carbon is normally present in the atmosphere in the form of gaseous compounds like carbon dioxide and methane. Carbon 14 C is a naturally-occurring radioisotope that is created from atmospheric 14 N nitrogen by the addition of a neutron and the loss of a proton, which is caused by cosmic rays.
This is a continuous process so more 14 C is always being created in the atmosphere. Once produced, the 14 C often combines with the oxygen in the atmosphere to form carbon dioxide. Carbon dioxide produced in this way diffuses in the atmosphere, is dissolved in the ocean, and is incorporated by plants via photosynthesis.
Animals eat the plants and, ultimately, the radiocarbon is distributed throughout the biosphere. In living organisms, the relative amount of 14 C in their body is approximately equal to the concentration of 14 C in the atmosphere. When an organism dies, it is no longer ingesting 14 C, so the ratio between 14 C and 12 C will decline as 14 C gradually decays back to 14 N.
This slow process, which is called beta decay, releases energy through the emission of electrons from the nucleus or positrons. After approximately 5, years, half of the starting concentration of 14 C will have been converted back to 14 N. This is referred to as its half-life, or the time it takes for half of the original concentration of an isotope to decay back to its more stable form.
Because the half-life of 14 C is long, it is used to date formerly-living objects such as old bones or wood. Comparing the ratio of the 14 C concentration found in an object to the amount of 14 C in the atmosphere, the amount of the isotope that has not yet decayed can be determined.
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